We thoroughly check each answer to a question to provide you with the most correct answers. Found a mistake? Tell us about it through the REPORT button at the bottom of the page. Ctrl+F (Cmd+F) will help you a lot when searching through such a large set of questions.

## CHEM 161 Lecture 1 Answers

What are the SI Base Units?
length (meter); Mass (kilogram); time (seconds); temperature (kelvin); amount of a substance (mole); electric current (ampere); luminous intensity (candela)

What’s the SI Base Unit for length?
meter

What’s the SI Base Unit for mass?
kilogram (kg)

What’s the SI Base Unit for time?
second (s)

What’s the SI Base Unit for amount?
mole (m)

What’s the SI Base Unit for temperature?
Kelvin

What’s the SI Base Unit for current?
ampere

What’s the SI Base Unit for luminous intensity?
candela

What is Planck’s constant?
6.626 x 10^-34 Js

What are the SI prefixes? What do they mean?
Tera (T) = 10^12 Giga (G) = 10^9 Mega (M) = 10^6 Kilo (k) = 10^3 Base Unit = 10^0 = 1 Milli (m) = 10^-3 Micro (u) = 10^-6 Nano (n) = 10^-9 Pico (p) = 10^-12

When are standard prefixes used for?
Standard prefixes are used for every three powers of 10

What is centi?
10^-2 c

Tera
T, 10^12, Trillion

Giga
G, 10^9, billion

Mega
M 10^6, million

Kilo
k, 10^3, thousand

Milli
m, 10^-3, 0.001

Micro
u, 10^-6, 0.000001

Nano
n, 10^-9, billionth

Pico
p, 10^-12, trillionth

Convert between Milli and Pico given 0.35 mm.
0.35 mm (1m/10-3 mm) (10^-12 p/1 m)

What are all the derived units?
Area: m², cm² Volume: m³, cm³ Force: N = Newton = kg(m/s²) = kgms⁻²) Energy: Joule = Nm = kg(m²/sec²) Pressure: P = Pascal = N/m² Density: kg/m³ g/cm³

How do you derive Force?
Force: N = Newton = kg(m/s²) = kgms⁻²)

How do you derive energy?
Joule = Nm = kg(m²/sec²)

How do you derive pressure?
P = Pascal = N/m²

How do you derive density?
kg/m³ g/cm³

What is a commonly used measure of volume?
liter

How are a mL and a cm3 related to each other?
they are equal 1000 mL = 1000 cm^3 = 1L

What’s the difference between mass and weight?
Weight: -equivalent to force (unit is Newtons) -a measure of gravitational attraction Mass: -amount of matter -a measure of a body’s inertia Inertia: resistance to change in motion (acceleration) -popcorn: easy to move back and forth -computer: hard to move back and forth

Why do we use significant figures?
Experimental measurements have an uncertainty. The uncertainty depends on the measuring instrument used

A length is measured as 4.98 cm. Assume the uncertainty is in the last digit (4.98 +- 0.01 cm). How many sig figs?
3

How many sig figs in 0.0498?
3 0.0498 +/- 0.0001 m

How many sig figs in 0.0000498 km?
3 0.0000498 +- 0.0000001 km

How can we prove that no matter how the measurements are converted, they have the same uncertainty?
Divide each of the numbers by their uncertainty. For each, the FRACTION UNCERTAINITY IS THE SAME = 1/500. 4.98/ 0.01 cm 0.0498 / 0.0001 m 30.0000498 / 0.0000001 km

Compare preceding zeroes and trailing zeroes. Which ones count as sig figs?
Only trailing zeroes or zeroes between other sig figs count as sig figs

What is the uncertainty of 48,000?
48,000 +/ 1000

How do you calculate the max/min of an uncertainty?
Using the original numbers, take the last decimal point of each of the numbers. For the min, move the number down 1 and for the max, move the number up 1. The difference between the Max and Min is the uncertainty.

How do you use sig figs when multiplying or dividing numbers?
Keep the same # of sig figs in your final result that is in the LEAST precise measurement of the calculation

How do you use sig figs when adding or subtracting numbers?
For adding or subtracting, keep the same number of DECIMAL PLACES in your final answer that is in the least precise measurement of the calculation

What are the exceptions to sig fig rules?
1. Exact Definitions (how many inches in a foot) 2. Counts of something (# cards in a deck) 3. Integral numbers that are part of an eq (bh/2) infinite # sfs

Convert 25 cm into m
25 cm * 1m/100cm = 0.25 m

Convert 10.0 cm into inches
10.0 cm * 1in/2.54 cm = 3.94 in (3 sf)

Convert 15.2 ft into inches
15.2 ft * 12 in/1 ft= 182.4 in= 182

What is the difference between an intensive and extensive property? What are examples of each?
Intensive properties, like density, don’t depend on the amount of matter. They are useful in describing and identifying a substance. Besides density, other intensive properties are MP and BP. Extensive properties would be mass or volume because they are different for each size of sample.

How do sfs dictate certainty?
The more sfs there are, the more precise your value is

How many sfs are there in 0.0540?
3 – trailing zeroes after a dc pt are sig

Why do we write numbers in scientific notation when we get large numbers (think sig figs)
trailing zeroes before an implied dc pt are AMBIGUOUS and scientific notation should be used instead (4000) 4 x10^3 = uncertainty +1000 (5000-3000 range) 4.0 x 10^3 = +(100) 4.00 x 10^3 = +10 4.000 x 10^3 = +1

(10.6-0.911)/(1.11)(0.3240)
First, 10.6 – 0.911 = 9.689 (2 sf) = 9.7 (don’t round yet tho in ur answer. just know there’s two sfs) Next, divide 9.689/ (1.11 3sf) (0.3440 4 sfs) = 26.94083 round to 2 sfs 27

What’s the difference between a physical and chemical property?
A physical property is something that can be observed or measured without changing the composition of matter

What is the difference between accuracy and precision?
Accuracy refers to how close the measured value is to the actual value. Precision refers to how close a series of measurements are to each other.

What kind of error persists with low accuracy and precision?
Random Error – error that has equal probability of being too high or too low. Usually averages out with repeated trials.

What kind of error persists with low accuracy but high precision?
Systematic error – error that tends toward being either too high or too low. However, unlike random error it doesn’t average out with repeated trials.

What is energy?
the ability to do work

What is work?
the action of a force through a distance

What is the total energy of an object?
Sum of Kinetic Energy + Potential Energy

What happens when you drop a weight in terms of energy?
It’s gravitational potential energy is converted into KE. When the weight hits the ground, its KE is converted primarily into thermal energy

If you drop a weight off a building, where is its KE highest? where is its PE highest? lowest?
On the tip of the building – high potential energy (unstable) Falling – highest KE On ground – low PE (Stable) + KE as heat

What is the law of conservation of energy?
Energy cannot be created or destroyed – the reactants = products

Why do objects with higher PE tend to be unstable?
Objects with higher PE try to change to lower their PE – makes them unstable

What is 1 J equal to?
1kg (m^2/s^2)

What is the formula of KE?
1/2mv^2 m = should be kg v = should be in m/s

What is an exothermic process?
when heat is released (lost) negative

What is an endothermic process?
a process that absorbs heat from the surroundings (gains) positive

18 gauge wire has a diameter of 1/25 of an inch. How many pounds of 18 gauge copper wire would have a length of 1.0 mile? Useful information: density of copper = 8.96 g/cm3, cost of bulk copper = \$3.03/lb, conductivity of copper=5.96x107A/Vm (at 20℃), Area of circle = πr2
ANS 26. lb

## Chem 161 Lecture 2 Answers

How can matter be classified?
Matter: Purity: 1. Pure Substances A. Elements B. Compounds 2. Mixtures A. Homogenous Solutions B. Heterogenous Solutions

What are the differences between compounds and mixtures?
Compounds: 1) Definite Composition (with exact chemical formula) 2) Have different properties than the elements making up the compound (H20 is liquid compared to H2 and O as gases) 3) Can be separated into its elements only by chemical reactions (water can only b separated by electrolysis NOT boiling) Mixture: 1) No definite composition (NaCl in H20 can be at various concentrations) 2) Has average properties of substances in the water ex. salty water 3) Can be separated into its components by physical means (filtration, distillation, chromatography)

You find a box filled with 35% sugar and 65% salt. What kind of substance is this?
Mixture – no definite composition. The individual proponents like sugar and salt have definite compositions (NaCl and C6H12O6) but in relation with each other, there is not necessarily the same composition found.

You find salty on your kitchen table. Is this a compound or mixture? Why?
It’s a mixture because you still get water and taste the salt. It would only be a compound if it had different properties compared to its elements. For example, for the water, both the H2 and 0 are gases by itself. however, together, it makes a liquid at room temp.

How do compounds and mixtures differ in being separating?
Compounds can be separated into its elements only by chemical reactions (water can only b separated by electrolysis NOT boiling) Mixtures can be separated into its components by physical means (filtration, distillation, chromatography)

How does water get separated back into H and O?
Electrolysis NOT boiling

What makes water a compound?
1) Water has the formula H2O, always has 2 atoms of hydrogen for one atom of oxygen. It has a definite composition by mass, always 88.9% oxygen, and 11.1 hydrogen. 2) Water is a liquid essential for life. Hydrogen is a light combustible gas; oxygen is a gas, which supports combustion. 3) Water can be broken down into hydrogen and oxygen by very high temperatures or electrolysis. These are chemical reactions.

Is common table salt, sodium chloride, a compound of mixture?
Compound: 1) NaCl always has 1 atom of sodium to one atom chloride. It has a definite composition by mass, always 39.3% sodium, 60.7% chloride. 2) Sodium chloride is a solid which dissolves in water and melts at 800 deg C. Sodium is a soft, low density metal which reacts violently with water. Chloride is a pale yellow green gas which is poisonous. 3) Sodium chloride can be broken down into sodium and chloride by electrolysis.

Is a solution of sodium chloride in water a mixture or compound?
MIXTURE: 1) A salt water solution has no definite composition. You can dissolve varying amounts of sodium chloride in a certain volume of water. 2) A salt water solution has the properties of both sodium chloride and water. 3) A salt water solution can be separated into its components by distillation, a physical process.

What were the four components of John Dalton’s postulates?
1) Each element is composed of tiny, indestructible particles called atoms A. partially true!! Yes, they are small particles. BUT, there are subatomic particles. 2. The same atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements A. Partially True. He was correct that each element has properties that are diff from other elements. BUT an atom of the same element isn’t necessarily the same mass. There are isotopes. 3. Atoms combine in simple, whole # ratios to form compounds A. COMPLETELY CORRECT 4. Atoms of 1 element do not change into atoms of another element. In a chemical reaction, atoms only change the way that they’re bound together to other atoms. A. Partially correct because TRUE, the number of atoms before and after are the same. BUT, atoms can change into another elements with nuclear reactions.

What is the first postulate of John Dalton’s atomic theory?
Each element is composed of tiny, indestructible particles called atoms

How true is Dalton’s first postulate of atomic theory?
Each element is composed of tiny, indestructible particles called atoms A. partially true!! Yes, they are small particles. BUT, there are subatomic particles.

What is the second postulate of John Dalton’s atomic theory?
The same atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements

How true is Dalton’s second postulate of atomic theory?
The same atoms of a given element have the same mass and other properties that distinguish them from the atoms of other elements A. Partially True. He was correct that each element has properties that are diff from other elements. BUT an atom of the same element isn’t necessarily the same mass. There are isotopes.

What is the third postulate of John Dalton’s atomic theory?
Atoms combine in simple, whole # ratios to form compounds

How true is Dalton’s third postulate of atomic theory?
Atoms combine in simple, whole # ratios to form compounds A. COMPLETELY CORRECT

What is the fourth postulate of John Dalton’s atomic theory?
Atoms of 1 element do not change into atoms of another element. In a chemical reaction, atoms only change the way that they’re bound together to other atoms.

How true is Dalton’s fourth postulate of atomic theory?
Partially correct because TRUE, the number of atoms before and after are the same. BUT, atoms can change into another elements with nuclear reactions.

What is Lavoisier’s Law of Conservation of Matter?
Mass of reactants = Mass of products

What law dictates that the Mass of reactants = Mass of products?
Lavoisier’s Law of Conservation of Matter Atoms combine or rearrange, but the total number of atoms of each type doesn’t change in a chemical reaction

What is Proust’s Law of Definite Composition?
Each compound has a definite percent of each element in that compound. This is sometimes called the Law of Definite Proportions. Ex. When broken down into hydrogen and oxygen, each 100 g of water produces 88.9 g of oxygen and 11.1 g of hydrogen

What is the law of definite proportions?
a chemical compound contains the same elements in exactly the same proportions by mass regardless of the size of the sample or source of the compound

What is the law of multiple proportions?
if two or more different compounds are composed of the same two elements, then the ratio of the masses of the second element combined with a certain mass of the first element is always a ratio of small whole numbers

X has 27.3 g C and 72.7 g O. Y has 42.9 g C and 57.1 g O. Does this support the law of multiple proportions? Why?
For X, divide each of the masses by 27.3 to get 1C: 2.66 O. You get a ratio of 1:3 as a result. For Y, divide each of the masses by 42.9 to get 1C and 1.33O. You get a 1:1 ratio. This supports the multiple proportion because X has 2x more O than Y. This would mean if C is CO2, Y is CO. Or if Y is CO2, then X is CO4.

How do cathode rays work?
Cathode ray tubes (late 19th century) had a low pressure gas and two electrodes under high voltage. A beam of light emanated from the negative electrode (cathode). This beam of light was deflected by both a magnetic and electric field, and the direction of the deflection showed that it consisted of negatively charged particles. These particles had the same properties regardless of the composition of the cathode.

How did JJ Thompson calculate the charge to mass ratio of an electron?
Charge of electron (e) / Mass of electron (m) = -1.758 x 10^8 coulombs/gram

Since Thomson found that all elements produced the same data, what does this indicate?
Thomson found that all elements produced the same data. That is, all elements produce this stream of negatively charged particles with the same charge to mass ratio, meaning that all elements contain electrons

What did Robert Millikan’s oil drop experiment do?
Millikan irradiated small sprayed oil drops with X-rays, which gave them a negative charge. These drops were then suspended between two charged plates By varying the charge on the plates, he could just suspend the oil drop, meaning the electrical force up (the negatively charged drop is attracted to the upper positive plate) is equal to the gravitational force down. Knowing the voltage on the plates, and through certain other measurements, Millikan could calculate the charge (in coulombs) on each oil drop. Doing this experiment many times, he came up with different charges on each drop. Some typical numbers are: 3.2 × 10-19 coulombs 6.4 × 10-19 coulombs 4.8 × 10-19 coulombs 9.6 × 10-19 coulombs 8.0 × 10-19 coulombs 11.2 × 10-19 coulombs This was done hundreds of time. Analysis shows that every measurement is an integer multiple of 1.6 × 10-19 coulombs. That is, each droplet of oil must have an integer number of excess electrons: 2 electrons, 4 electrons, 3 electrons, etc.

What is the charge of one electron according to Millikan?
e = 1.602 x 10^-19 coulombs

Combining Millikan’s experiment with Thompson, what did we find?
We found the mass of an electron. Millikan: e = 1.602 x 10^-19 coulombs JJ. Thomson: e/m = -1.758 x 10^8 coulombs/gram m = mass of electron = 9.11 x 10^-28 g

What is the mass of an electron?
9.11 x 10^-28 g

What was Rutherford’s experiment?
the gold foil experiment – positively charged alpha particles were fired at a very thin sheet of gold foil. Most of the alpha particles went through undeflected BUT some were deflected at very large angles. They deflected because they hit or come close to the nucleus. But because only a few were deflected, it showed that the nucleus must be very small (but most of the mass) and that the electron cloud is very big but has a smaller mass

What was the conclusion of the Rutherford experiment?
Atoms contain a nucleus at the center which is relatively very heavy (when the alpha particles hit it they bounce back at you) and also very small (most of the alpha particles don’t hit it at all and go through the atom undeflected) through the electron cloud.

What did Rutherford’s Scattering Experiment disprove?
Plum Pudding Model

What is the mass (AMU) of a proton?
1.00727

What is the mass (g) of a proton?
1.673 x 10^-24

What is the mass (AMU) of a neutron ?
1.00866

What is the mass (g) of a neutron?
1.675 x 10^-24

What is the mass (AMU) of a electron?
0.0005486

What is the mass (g) of a electron?
9.11 x 10^-28

What is the ratio of the mass of the proton compared to the mass of the electron? ?
1.00727 / 0.0005486 =1836
How do you depict an element with atomic number and mass number?
Mass number on top of Atomic number

How do you calculate average atomic mass?
For all naturally occurring isotopes: (%abundance x mass) + (%abundance x mass) + etc…

How do you calculate the percent abundance of each isotope?
Cu-63 Mass: 62.9298 %abu: ? Cu-65 Mass: 64.9278 %abu: ? First, set the %abu of Cu-63 to x and the %abu of Cu-65 to 1-x. The sum of the two percentages must be equal to 100% basically, its easier to express as a decimal. You know the average mass from the periodic table. Set the weighted average mass = 62.9298 x + 64.9278 (1-x) = 63.546 Solving: x = 0.6916 or 69.16% 1-x = 0.3084 or 30.84%

What is deuterium?
1 proton and 1 neutron Given symbol D

Compare D2O to H20
Water w deuterium has a higher density than liquid H20. Ice with D2O sinks.

How did Millikan find out the charge of the electron?
If you are given the charge of an oil drop, divide it by charge of another or same oil drop that is the smallest among the oil drops. You should get a whole number ratio. If not, make the denominator 1 by multiplying. Afterwards, take the initial charge you were given and divide it by the number you got from dividing the two charges (number of electrons). As a result, you should get the fundamental charge of an electron

What is the fundamental charge of an electron?
-1.60 x 10^-19 C

What is the label for AMU?
amu/atom or g/mol

## Chem 161 Lecture 3 Answers

What is a mole?
number of atoms that has a mass equal to the AMU in grams

6.022 x 10^23 atoms/mole
What is wavelength?
distance between each wave unit λ

What is frequency?
# of waves (cycles) per second

What are units of wavelength?
m, mm, nm (length of each λ)

What are units of v?
cycles/sec waves/sec sec^-1 Hz (heartz)

What is the velocity of the wave?
λ * v = velocity of wave (meter/wave) * (wave/sec) = meter/sec

What is the speed of light?
2.998 x 10^8 m/s = c

For electromagnetic radiation, c = ?
λv=c

Electromagnetic radiation encompasses what kinds of rays?
radio waves, infrared, visible light, ultraviolet, x-rays, and gamma rays

Visible light is a small portion of the spectrum with wavelengths ranging from ?
400-700 m

What does c designate?
speed of light in a vacuum c = 2.998 x 10^8 m/s

What is the relationship between frequency and wavelength?
inverse

Order the electromagnetic spectrum in terms of increasing frequency and increasing wavelength
—> Inc frequency <– Dec frequency

What is the relationship between energy and frequency and wavelength?
Energy and frequency are proportional Energy and wavelength are inverse

Order the electromagnetic spectrum

Order the electromagnetic spectrum in increasing frequency

Order the electromagnetic spectrum in increasing wavelength
gamma rays –> UV –> visible –> infrared –> radio

With visible light, which colors have the highest frequency?
Violet has highest frequency Red has lowest frequency

With visible light, which colors have the highest wavelength?
Red has highest wavelength Violet has lowest wavelength

Order visible light in increasing frequency
red, orange, yellow, green blue, violet

Order visible light in increasing wavelength
violet, blue, green, yellow, orange, red

What was the wavelength of red? What was the frequency?
l = 700 nm v = 4.3 x 10^14 s^-1

What was the wavelength of violet? What was the frequency?
l = 400 nm v = 7.5 x 10^14 s^-1

What impacts the color of the radiation?
The color of the radiation is determined by the frequency ( and wavelength in association with lv = c)

What determines the brightness or intensity of the radiation?
The amplitude (height) of the wave

How does frequency and radio waves relate?
A typical FM radio station is about 100 MHz or 10^8 s^-1. Visible light has frequency in the range of 10^14 sec^-1 or about a million times higher than radio waves.

What is the wavelength of the radio station with frequency 100 MHz?
100 MHz is the same as 100 x 10^6 or 1 x 10^9 Hz or 1 x 10^8 s^-1 l = c/v = (2.998 x 10^8 m/s) / (1 x 10^8 s^-1) = 3

The wavelength of radio waves is much BLANK than that of visible light
longer (3 m vs. 400-700 nm)

Compare the frequency of visible light to the frequency of radio waves
The frequency of visible light is much higher than the frequency of radio waves

Compare the wavelength of visible light to the wavelength of radio waves
The wavelength of visible light is much higher than the wavelength of radio waves

wave-particle duality of light
Light sometimes acts like a particle and sometimes like a wave

How are the electric field and magnetic field component of light related to each other?
90 degrees to each other

What is diffraction?
Important wave behavior: bending of light around an obstacle or through a slit the aperture effectively becomes a secondary source of the propagation of wave

What is interference?
Important wave behavior: separate light waves can interact by overlapping and either building up or cancelling each other

What is constructive interference?
Where two waves arrive in step reinforcing one another (increasing the amplitude) in phase

What is deconstructive interference?
the trough of one wave overlaps the crest – evens out out of phase

What are photons?
particles/quanta of light

Energy of a photon = ?
Planck’s constant x frequency = J (Planck’s constant * speed of light)/wavelength

What is 1 Watt equal to?
1 joule per second

A 25 Watt red light bulb emits a wavelength of 667 nm.
Energy of photon = hc/l =(6.626×10&-34 Js)(2.998×10^8m/s) / (6.67 x 10^-7 m) Ephoton = 2.98 x 10^-19 J/photon l = 667 nm (1 m/10^9 nm) = 6.67 x 10^-7 m 1.0 hour * (60 min/1 hr) * (60 s/1 min) * (25 J/1 s) * (1 photon/2.98 x 10^-19 J) = (3.0 x 10^23 photons)

How can the photoelectric effect of light be explained?
Particle nature of light Shine a radiation source on a metal surface and you see light (electrons) being emitted from the metal surface

Experimental Findings of Photoelectric Effect
1. For a given metal, there is a characteristic minimum frequency (v0) of light needed for the photoelectric effect to occur 2. If light is below threshold frequency, increasing duration of irradiation of light intensity has NO EFFECT 3. Once the threshold frequency is met/exceeded, increasing the intensity causes an increase in the NUMBER of photoelectrons emitted, but not their velocity 4. Increasing the frequency beyond the threshold frequency increases the velocity of photoelectric effect

What is minimum frequency (v0)?
the threshold frequency

Why can’t the wave nature of light explain the photoelectric effect?
1. If light is acting as a wave, increased duration and intensity should eventually lead to the photoelectric effect, even if it takes a long time 2. If light is acting as a wave, increasing intensity should also lead to faster electrons (but not necessarily more electrons) being ejected from the metal

Photoelectric Effect Equation
Ephoton (hc/l) = Binging energy (BE) hv0 (threshold frequency) + KE (1/2mv^2) m = kg v = m/s

Binding energy is the same thing as?
threshold energy and work function ∅

Threshold energy ∅ is not the same as
threshold frequency (v0)

If Ephoton < BE?
no photoelectric effect

If Ephoton > BE
BE met and leftover will be used as KE

If Ephoton >> BE
photoelectric effect with higher velocity and greater KE

## Chem 161 Lecture 4

What is the threshold frequency?
v0 = if you vary the frequency of the incident light, it is found that there is a certain minimum threshold frequency below which no electrons are emitted

What happens if you are below the threshold frequency?
No electrons are emitted regardless of the brightness of the light

What happens if you are above the threshold frequency?
Electrons are emitted, and the higher the frequency, the higher the KE of the emitted electrons

What is the formula for binding energy + KE?
hv (energy of the incoming photon) = hv (binding energy of electron; v0 = threshold frequency) + KE (KE of emitted electron) rsm

Compare KE and slope
KE = hv – hv0 y = mx + b

What does the intensity (brightness) of the light correspond to?
# of photons

What does a brighter light indicate about photons?
There are more photons

How do photons and electrons interact?
A single photon can be absorbed/emitted by a single electron which gains/loses the energy of the photon. 1:1 (one photon in: one photon out) hv>hv0 – need sufficient energy

What happens if hv < hv0?
each photon doesn’t have enough energy to ionize the atoms of the metal, regardless of the number of photons (intensity) of the incident light

What happens if frequency is too low?
no photons are emitted

What happens with light of high-enough frequency?
a brighter light results in more photons emitted, but the energy of each of those photons remains the same

When light of frequency 7.00 × 1014 sec-1 strike an emitter, electrons with kinetic energy 2.00 × 10-19 J are emitted. What is the threshold frequency?
(6.626E-34)(7.00E14) = hv0 + 2.00E-19 hv0 = 2.638E19J v0 = (2.638E-19)/(6.626E-34) = 3.98E14 s-

In a photographic darkroom, why is the safelight red?
Red is the lowest frequency and lowest energy photon of visible light. Red light photons are not energetic enough to expose black & white photographic paper.

A person has a roll of photographic film in a completely dark room with an X-ray machine on. When developed the roll of film is completely exposed (black). Why?
X-rays are high-energy photons which are even more energetic than visible light. These photons expose film (activate the silver salts in the film) and when developed, black silver forms everywhere

A person has a roll of photographic film and visits a powerful radio station transmitter. When developed, the film is normal. Why?
The radio station has lots of photons (high intensity), but the photons have low frequency and hence low energy. They are not energetic enough to expose the film.

The energy of each photon is determined by the BLANK. The intensity (brightness) of the light is determined by the BLANK.
The energy of each photon is determined by the frequency of light: E =hν. The intensity (brightness) of the light is determined by the number of photons.

How many photons are present in a 5.0 milliwatt red laser pointer, λ = 635 nm, when on for 6.0 seconds.
Watt = unit of power = Joule/sec Total energy of laser pointer when on for 6.0 seconds: (5.0×10^-3 J/sec) * (6sec) = 0.030 J Energy of each photon = hv = hc/l. For the laser pointer, l = 635 nm E = 3.13E-19J Number of photons needed for 0.030J total energy: 0.030 J * (1photon/3.13E-19) = 9.6E16

Binding energy of an electron = 193 kJ/mol What is the threshold frequency
E = hv Common mistake: 193 = 6.626 × 10-34 (ν) Solve for ν Why is this wrong? Units must be consistent. The binding energy is given in kJ/mol. The value of h is in J s, and is not per mole. Correct sol = E = hv 193000 J/mol * 1mol/6.022xE23 =6.626E-34 Js (v) Solve for v ν = 4.84 × 1014 s-1

What is the emission spectrum of an excited atom?
The emission spectrum from excited atoms looks different. When examined with a prism, there is no continuity of colors, but rather a series of distinct lines of varying wavelengths and colors. Each atom has its own particular pattern (emission spectrum).

What is the expression for energy levels?
E = -2.178 x 10^-18 J (Z^2/n^2) For hydrogen, Z = 1 = -2.178 x 10^-18 / n^2 J

Each line of the hydrogen atom spectrum results from an electron changing levels. In an emission spectrum, the excited electron starts at a higher level and goes down to a lower level, emitting light of a specific frequency and wavelength. The frequency can be calculated by:
Delta E = hv where E is the difference in the two levels

What does each line in the spectrum correspond to?
Each line in the spectrum corresponds to a “jump” between levels, a transition from a higher level to a lower, emitting a specific quantity of energy, and a specific frequency and wavelength. The

What can the frequencies and wavelengths of each line of an emission spectra be calculated by?
delta E = hv = hc/l = Ef-Ei = 2.178×10^-18 (1/ni^2 – 1/ni^2)

For emission spectra, we start at a higher level and go down to a lower, thus doing what to energy
emitting

For emission spectra, we start at a higher level and go down to a lower, thus emitting energy. This means what about n?
This means ni > nf, which logically means that 1/ni^2 > 1/nf^2 meaning that ΔE will be negative. This makes sense since the final energy is lower than the initial energy.

What is the energy emitted when the electron goes from n = 3 to n = 2?
Delta E = 2.178 x10^-18 (1/3^2 – 1/2^2) = -3.025 × 10^-19 Joule Note that ΔE is negative since energy is released. This will always occur if ni > nf, since we are starting at a higher level and going down to a lower level.

What frequency and wavelength of light emitted when the electron goes from n = 3 to n = 2?
Since ν and λ must be positive, we’ll use a positive value for ΔE. ΔΕ = energy released = 3.025 × 10^-19 Joule v = delta E/h = (3.025×10^-19J) / (6.626×10^-34Js) = 4.565 × 1014 s-1 l = c/v = (2.998×10^8 ms^-1)/(4.565×10^14 x^-1) = 6.567 x10^-7 m = 656.7 nm

When we calculate v or l from emission spectra delta E, we use what of the E?
Note that when we calculated ν or λ, we used the absolute value of ΔE. ΔE is negative when energy is released. But a negative value of ν or λ would not make sense. The negative ΔE means that energy is released, and light of a particular frequency or wavelength is emitted.

What is the Balmer series?
transitions from n = 2, energies occur in the visible region

What is the Lyman series?
any transition into or out of the ground state (n=1) (all in the UV)

What is the Paschen series?
series of lines with nf = 3 is called the Paschen series.

How much energy is required to remove the electron out of the atom completely, starting with the electron in the ground state (n = 1)?
When out of the atom, n = ∞, and E = 0 Delta E = 2.178 x10^-18 (1/ni^2 – 1/nf^2) =2.178 x10^-18 (1/1^2 – 1/infinity^2) =2.178 x10^-18 J

What is the E and n of an electron out of the atom?
When out of the atom, n = ∞, and E = 0

What is the ionization energy of hydrogen?
This is known as the ionization energy of hydrogen–the energy needed to remove the electron from the atom.

What is often expressed in units of kJ/mol?
ionization energy of hydrogen 2.178 x10^-18 J/atom * 6.022×10^23 atom/mol *1 kj/1000J = 1312 kJ/mol

What is the DeBroglie Wavelength?
lambda = h/mv = h/mc E = mc^2 for any particle of mass v and velocity v

What is the wavelength of an electron which moves at 1/100 of the speed of light?
For an electron, m = 9.11 × 10^-31 kg For this electron, v = 3.00 × 10^6 m /sec l = h/mv = (6.626E-34)/(9.11E-31 kg)(3.00E6 m/s) = 2.42E-10 m = X – ray wavelength

What is the Bohr constant?
B = 2.179 x 10^-18 J

How to find the energy of electron in energy level?
En = -B/n^2 B = 2.179 x 10^-18 J n = 1,2,3,4 E1 = -2.179×10^-18J (ground state) Excited States; E2 = -5.447 x 10^-19 J E3 = -2.421 x 10^-19 J E4 = -1.362 c 10^-19 J

How can the electron be promoted to a higher energy level?
by absorbing energy (as a photon of light)

How can the electron drop down to a lower energy level?
by emitting energy (as a photon of light)

How does the energy of the photon mathematically relate to electronic transitions?
Ephoton (hv = hc/lambda) = -B(1/ni^2- 1/nf^2)

If the sign of delta E or energy of the photon is positive, then?
photon absorbed energy nf>ni

If the sign of delta E or energy of the photon is negative, then?
photon emitted energy ni>nf

What is the limitation of the Bohr equation for line-spectra?
can only predict one electron systems H, He+, Li2+

What unit is the de Broglie wavelength given in?
m

What unit is the de Broglie wavelength’s mass given in?
kg – mass of particle

What unit is the de Broglie wavelength’s velocity of particle given in?
m/s

## Chem 161 Lecture 5 Answers

What are the four quantum values?
n is the principal quantum number. It determines the main energy level. Values of n are 1,2,3,4,5,…… ℓ is the angular momentum quantum number. It determines the shape of the orbital Values of ℓ vary from 0 to n-1 mℓ is the magnetic quantum number. It determines the orientation of the orbital. Values of mℓ vary from -ℓ to +ℓ

What is n?
principal quantum number

What does n determine?
main energy level.

What are the values of n?
Values of n are 1,2,3,4,5

What is l?
angular momentum quantum number

What does l determine?
It determines the shape of the orbital

What are the values of l?
Values of ℓ vary from 0 to n-1

What is mℓ?
magnetic quantum number

What does mℓ determine?
It determines the orientation of the orbital.

What are the values of m?
Values of mℓ vary from -ℓ to +ℓ

Whare are the l, ml, and orbital designations of n=1?
l = 0 ml = 0 orbital designation: 1s

Whare are the l, ml, and orbital designations of n=2?
n=2: l = 0 ml = 0 Orbital designation: 2s n=2: l = 1 ml = -1,0,+1 orbital designation: 2p

Whare are the l, ml, and orbital designations of n=3?
l = 0 ml = 0 orbital designation: 3s l=1 ml= -1,0,+1 orbital designation: 3p l=2 ml = -2,-1,0,+1,+2 orbital designation: 3d

Whare are the l, ml, and orbital designations of n=4?
l = 0 ml = 0 orbital designation: 4s l = 1 ml = -1,0,+1 orbital designation: 4p l = 2 ml = -2,-1,0,+1,+2 orbital designation: 4d l = 3 ml = -3,-2,-1,0,+1,+2,+3 orbital designation: 4f

How many total orbitals are there in n=1?
For the first energy level, (n=1) there is one type of orbital (1s) and a total of one orbital (1s)

How many total orbitals are there in n=2?
For the second energy level (n=2), there are two types of orbitals (2s,2p) for a total of four orbitals (2s + three orientations of 2p, designated as 2px, 2py, 2pz). All p orbitals (ℓ = 1) have three orientations (mℓ = -1,0,+1).

How many total orbitals are there in n=3?
For the third energy level (n=3), there are three types of orbitals (3s,3p,3d) for a total of nine orbitals (3s + three 3p orbitals + five 3d orbitals). All d orbitals (ℓ = 2) have five orientations (mℓ = -2,-1,0,+1,+2)

How many total orbitals are there in n=4?
For the fourth energy level (n=4), there are four types of orbitals (4s,4p,4d,4f) for a total of sixteen orbitals (4s + three 4p orbitals + five 4d orbitals + seven 4f orbitals). All f orbitals (ℓ = 3) have seven orientations (mℓ = -3,-2,-1,0,+1,+2,+3)

What is the mathematical relationship between energy level, orbital, and total orbital numbers?
For the nth energy level, there are n types of orbitals for a total n^2 orbitals.

What is a node?
particular distance where the probability is zero. There is no easy classical explanation for the node; it is a consequence of the wave nature of the electron.

Why is the 2p orbital not spherically symmetric?
It is oriented along an axis, either x, y, z. There are three orientations of the 2p orbital, called 2px, 2py, and 2pz, each one pointing along one of the three mutually perpendicular axes.

How many potential orientations are there for s?
1 = 2e

How many potential orientations are there for p?
3 = 6e

How many potential orientations are there for d?
5 = 10e

How many potential orientations are there for f?
7 = 14e

Order the energy levels in increasing energy
1s 2s 2p 3s 3p 4s 3d 4p 5s 4d 5p 6s 4f 5d 6p 7s 5f 6d 7p 5g 6f 7d

What are the two principles that are used to determine electron configuration?
Aufbau Principle “Aufbau” literally means “build up” in German. This means we first place electrons in the orbitals with the lowest energy continue working our way up to the higher energy levels. Pauli Exclusion Principle This principle states: No two electrons in an atom can have the same set of quantum numbers.

What is the fourth quantum number?
This fourth quantum number is called the spin quantum number, ms, which can have values of either +½ or -½ . This is related to the magnetism inherent in electrons: the electron can have a magnetic moment with two possible values. It can be interpreted as the electron spinning in two opposite directions, although this is not a literal description
What is Hund’s rule?
When filling orbitals of equal energy, we keep electrons in separate orientations with parallel spins, pairing only when necessary.

When does the 4f sublevel begin?
after 6s the lanthanides are filling the 4f sublevel, then 5d transition metals

what is the electron configuration of actinides?
5f

Chromium electron config
[Ar] 4s1 3d5 The 4s orbital as well as each of the five 3d orbitals are singly occupied, for a total of 6 unpaired electrons.

Copper electron configuration
[Ar] 4s1 3d10

Electron Configurations of Anions
Negative ions (anions) are formed by adding electrons to the atom. The electrons are added to the lowest-energy empty orbital available. This generally results in a configuration identical to that of a noble gas.

Cl -Cl
Cl 1s2 2s2 2p6 3s2 3p5 Cl- s2 2s2 2p6 3s2 3p6

Electron Configurations of Cations
Positive ions (cations) are formed by removing electrons from the highest energy orbital. This is generally the orbital with the highest principal quantum number. For main group elements, this will usually result in an ion with the same electron configuration as a noble gas.

Na Na+
Na 1s22s22p63s1 Na+ 1s22s22p6

What happens when transition metals form positive ions?
When transition metals form positive ions, the electrons are first removed from the orbital with the highest principal quantum number, which, for the first transition series, would be 4s, not 3d.

Fe Fe+2 Fe+3
1s2 2s2 2p6 3s2 3p6 4s2 3d6 1s2 2s2 2p6 3s2 3p6 3d6 1s2 2s2 2p6 3s2 3p6 3d5

Zn Zn+2
1s2 2s2 2p6 3s2 3p6 4s2 3d10 1s2 2s2 2p6 3s2 3p6 3d10

Tim and lead are main group elements (4A) which usually form what kind of ions?
forms +2 ions

Tin and lead are main group elements (group 4A), which usually forms 2+ ions. The electrons being removed come from what energy level?
Tin and lead are main group elements (group 4A), which usually forms 2+ ions. The electrons being removed come from the highest-energy level, in this case 4p or 5p, with the resultant ion not being the same as that of a noble gas.

Sn Sn+2
Sn [Kr]4d105s25p2 Sn2+ [Kr]4d105s2

Pb Pb+2
Pb [Xe]4f145d106s26p2 Pb2+ [Xe]4f145d106s2

paramagnetic vs diamagnetic
paramagnetic- contains at least one unpaired electron and is referred to as radicals diamagnetic- contains no unpaired electrons

What is the de Broglie wavelength associated with the electron removed from tungsten (work function of 4.32-eV) when bombarded with ultraviolet light of 282.-nm?

Consider three metal samples designated A, B, C. Metal A emits electrons in response to visible light while metals B and C require UV light. What is metal A and what is the longest wavelength that can eject an electron from it? What range of wavelengths would distinguish metal B from C (i.e. cause only one metal to eject an electron)? MetalDensity (g/cm3)Work Function (J)Melting Temp. (K)Ba3.5104.30×10-19 J1000. KTa16.654.81×10-19 J3290. KW19.257.16×10-19 J3695. K

What is ionization energy?
Ionization energy is defined as the energy required to remove an electron from a gaseous atom.

What is the formula for ionization energy?
X (g) + energy –> E+(g) + e-

What are the periodic trends for IE?
Horizontal: IE increases across a row of the periodic table (from left to right) Vertical: IE decreases down a group (column) of the periodic table (from top to bottom)

As you go left to right horizontally, what happens to IE?
increases

As you go down vertically, what happens to IE?
decreases

Why does IE increase as you move left to right horizontally?
As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher Z, higher effective nuclear charge) attracts the electrons more, making it harder to remove an electron from the atom–hence a higher IE.

Why does IE increase as you move down vertically?
As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. It is thus easier to remove electrons from outer shells, which are less attracted to the nucleus

What is the exception to where ionization energy decreases from left to right?
Every element in group 3A (B, Al, Ga, In, Tl) has an ionization energy that is lower than the previous element.

Why does every group in 3A have a lower ionization energy compared to the previous element?
Elements in Group 3A all start the p block, with outer shell p1. The p electrons are somewhat higher in energy than s, since the p electrons are more effectively shielded from the nuclear pull (s electrons penetrate the core). When we start the p block, we’re removing the last electron from the p orbital, and this is a little easier to do than the previous element, giving rise to a lower ionization energy for this group.

Other than elements in 3A, what is the exception to the general trend of IE?
has a lower ionization energy than N; similarly, S has a lower ionization energy than P. This is related to Hund’s rule, which places electrons in separate orientations for the lowest possible energy. Nitrogen has a half-filled sublevel: each of the three p electrons are in separate orientations. The oxygen configuration has one of the p sublevels with paired electrons. The last electron in O must start pairing on the p sublevels, which results in a certain amount of repulsion, being in the same region of space. This results in a slightly lower ionization energy.

What is the second ionization energy’s definition and formula?
The second ionization energy, I2, is the energy needed to remove the second electron; I3 is the energy needed to remove the third electron, and so on. X+(g) + I2 → X2+(g) + e- X2+(g) + I3 → X3+(g) + e-

Why does ionization energy increase as you move to move the next electron in the atom?
proton. This ion holds the remaining electrons more tightly, making the second electron harder to remove. If the second electron is removed, we now have a 2+ ion, an excess of 2 protons, which holds the remaining electrons even more tightly. If the third electron is removed, we now have a 3+ in which in turn holds the remaining electrons even more tightly.

When does the big jump occur in iE?
The “big jump” will occur at different ionization energies depending on the number of valence electrons

Ne, Na+, Mg2+, Al3+ are all isoelectronic. They all have the same 10 electrons, with configuration 1s22s22p6. Which would have the highest ionization energy?
The reason: although all four isoelectronic species have the same 10 electrons, they have different number of protons. Ne, Na+, Mg2+, Al3+ have 10, 11, 12, 13 protons respectively. For the same number of electrons, the one with the most protons will hold onto those electrons most tightly and consequently will have the highest ionization energy.

Atomic radii measure the size of the atom. Although the atom does not have a distinct boundary, there are several ways to estimate atomic radii based on distances between atoms in crystals or molecules.

What is the periodic trend as you move from left to right horizontally for atomic radii?
atomic radius decreases across a row of the periodic table (from left to right)

What is the periodic trend as you move from down vertically for atomic radii?
atomic radius increases down a group of the periodic table (from top to bottom)

Why does atomic radii shrink as you move left to right horizontally?
As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher Z) attracts the electrons more, making it the electron cloud closer to the nucleus; hence a smaller atomic radius.

Why does atomic radii shrink as you moved down vertically?
As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. So the outer shell logically is larger there are more shells (higher principal quantum number).

What is the relationship between IE and Atomic Radii trends?
Note that the atomic radii trends are just the reverse of those for IE, but with the same reason. Going from left to right across a row, the IE generally increases, and the atomic radius decreases, both due to increased nuclear charge as the shell is being filled. going from top to bottom down a column, the IE decreases and the atomic radius increases, both due to more shells which are further from the nucleus and shielded from the nuclear pull by the core electrons.

Positive ions are formed when the atom gives off an electron. The resultant ion is always smaller than the corresponding atom, since the resultant positive charge (more protons than electrons) causes the remaining electron cloud to be pulled toward the nucleus.

Negative ions are formed when the atom gains an electron. The resultant ion is always larger than the corresponding atom, since the resultant negative charge (more electrons than protons) causes the electron cloud to be repelled away from the nucleus.

What is electron affinity?
When an electron is added to a gaseous atom, forming a negative ion, energy may be either released or absorbed. The electron affinity is defined as the change in energy.

When energy is released, the electron affinity is ?
negative

The more negative the electron affinity, the more the?
attraction.

Sometimes, electron affinity is positive, meaning that?
must be added in order for the electron to add on to the atom.

What is equal to energy change?
electron affinity X(g) + e- → X- (g) Εnergy change = electron affinity

If electron affinity is negative, energy is
released

If electron affinity is positive, energy is
absorbed (Required)

What are the trends in electron affinity?
Horizontal: Going from left to right across a row, the electron affinity gets more negative (more attraction for electrons), with the halogens (not the noble gases) having the most negative electron affinity (most attraction for electrons). Vertical: Going from top to bottom down a group, the electron affinity gets less negative (less attraction for electrons).

Why does electron affinity increase as you go left to right horizontally?
As you go across from left to right in a row of the periodic table, you are adding electrons to the same shell but with increasing nuclear charge. The increasing number of protons (higher effective nuclear charge) attracts the electrons more; a higher attraction for electrons means a more negative electron affinity

Why does electron affinity decrease as you move down vertically?
As you go down a group from top to bottom, you always have the same valence shell configuration. However, each succeeding shell is further from the nucleus, and is shielded from the nuclear pull by inner electrons. Less attraction by the nucleus means a less negative electron affinity.

What are the exceptions to electronic affinity?
to the nucleus. Similarly, oxygen has a less negative electron affinity than sulfur, despite having its outer shell closer to the nucleus. The reason for this is related to the extra electron repulsions in the small fluorine and oxygen atoms

What are the electron affinities for groups 2A and 8A?
Groups 2A and 8A have positive electron affinities, meaning that energy must be added to form negative ions of these elements. For group 2A, any added electron would have to go into the next p orbital, which, being more shielded from the nuclear pull than s orbitals, would have less attraction for incoming electrons. For the noble gas group 8A, added electrons would have to go into the next highest s orbital, which is highly shielded from the nuclear pull by the inner shells.

Why is electron affinity often positive the second time you add an electron?
Adding a second electron to the gaseous atom, producing a -2 ion, always requires energy, due to the repulsion of that electron from the previous -1 ion formed. This is the case even for something with a high negative electron affinity, such as oxygen. The high positive number for EA2, +744 kJ, means that 744 kJ must be added to place a second additional electron on a gaseous oxygen atom, forming a -2 oxide ion.

Why are the noble gases so unreactive?
They have a very high ionization energy–therefore it is difficult to remove electrons. They also have a positive electron affinity–it requires energy to add an electron since their p orbitals are complete, and any added electrons would have to go into an outer shell which is shielded from nuclear attraction. This combination of high ionization energy (hard to remove electrons) and positive electron affinity (difficult to gain electrons) makes the noble gases very unreactive.

Metallic Character
In general, the ability to act as a metal decreases from as you go from left to right across a row of the periodic table. This is because, as you go across the row, you add one more electron to the same shell, but with increasing number of protons, the electrons are held more tightly, resulting in higher ionization energy, smaller atomic radius, as well as decreasing metallic character. In general, the ability to act as a metal increases as you down a group of the periodic table. As you go down the group, the outer electrons are further from the nucleus, resulting in a lower ionization energy, with the outer electrons being held less tightly. This correlates with greater metallic character.

## Chem 161 Lecture 7 Answers

What is ionic lattice?
An ionic lattice is a long-range array of positive and negative ions, held together by electrostatic attraction

What is ionic lattice measured in?
kJ/mol

What does lattice energy depend on?
These comparisons can be answered by recognizing that lattice energy depends on both the charge of the ion and the size of each ion—which determines the distance between the oppositely charged ions

What formula shows the force of attraction between the positive and negative ions, and the energy needed to break this force of attraction?
Force / Energy = kq₁q₂/r² r stands for the distance between ions q stands for the charge on each ion

What happens to lattice energy as atomic radii increases? Would you expect KI to have a higher or lower lattice energy than NaCl?
“r” stands for the distance between the ions. When the ions are larger (K+ is larger than Na+ and I- is larger than Cl-), the interionic distance “r” is greater, leading to weaker electrostatic attraction and lower lattice energy KI: -632 kj/mol NaCl: -787 kj/mol

What happens to lattice energy as charges increases? Would you expect MgO to have a higher or lower lattice energy than NaCl?
The lattice energy depends on the force of attraction between the positive and negative ions, and the energy needed to break this force of attraction. “q” represents the charge on each ion. MgO has electrostatic attraction between +2 and – 2 ions which should be 4 times greater than the +1 and -1 ions in NaCl, hence the much larger lattice energy.

How much more energy is required to break apart an ionic bond between a +2/-2 ion compared to a +1/-1 ion
Energy = kq1q2/r 4 times more

How much more energy is required to break apart an ionic bond between a +1/-1 ion if the distance between the ions is halved
Energy = kq1q2/r times more. 2 times more. r (distance between charges) in the denominator is halved– so the electrostatic energy is doubled.

If the charge on one ion is doubled and the distance between the ions is also doubled, what effect does this have on the energy of the bond?
Energy = kq1q2/r times more. no change in energy. One of the charges in the numerator double, and r in the denominator doubles. The doubling cancels out.

What kind of ions do Group 3A metals form?
Group 3A metals, most commonly aluminum, form 3+ ions

What kind of ions do Group 7A halogens form?
All group 7A halogens form 1- ions

What kind of ions do Group 6A elements form?
All group 6A elements form 2- ions

What kind of ions do Group 5A metals form?
Group 5A elements, most commonly nitrogen, phosphorus form 3- ions

Suffix: “ate”
contains 3 or 4 oxygens + the element that’s in the name

Suffix: “ite”
contains one less oxygen than the corresponding “ate” ion but with the same charge

Suffix: “bi_ate” or “hydrogen_ate”
add one hydrogen atom to the corresponding “ate” ion and increase the charge by 1.

Suffix: “bi_ite” or “hydrogen_ite”
add one hydrogen atom to the corresponding “ite” ion and increase the charge by 1

SO4 2-
Sulfate

SO3 -2
Sulfite

NO3-
Nitrate

NO2-
Nitrite

ClO4-
Perchlorate

ClO 3-
Chlorate

ClO-2
Chlorite

ClO-
Hypochlorite

CO3 2-
Carbonate

HCO3-
Bicarbonate (Hydrogen Carbonate)

PO4-3
Phosphate

CrO4 2-
Chromate

HSO4-
Bisulfate or hydrogen sulfate

HPO4 -2
Hydrogen phosphate

H2PO4 –
Dihydrogen phosphate

O2-
oxide

OH-
Hydroxide

NH4+
Ammonium

Sodium sulfate
Na2SO4

Magnesium Sulfate
MgSO4

Aluminum sulfate
Al2(SO4)3

Sodium carbonate
Na2CO3

Sodium bicarbonate
NaHCO3

Iron (III) nitrate
Fe(NO3)3

Ammonium chloride
NH4Cl

ammonium sulfate
(NH4)2SO4

Diatomic Hg ion
Hg2 +2

mercury (1) chloride
HgCl2

Mercury (11) chloride
HgCl2

Copper (II) sulfate pentahydrate
CuSO4 * 5H2O

Barium chloride dihydrate
BaCl2 * 2H2O

Magnesium sulfate heptahydrate
MgSO4 * 7H2O

hydrate vs. anhydrate
Hydrate: CoCl2 * 6H20 Anhydrate: CoCl2

How do we represent the covalent bond?
We represent the covalent bond with a line between the two atoms, the line representing two electrons, one from each atom.
Polar Covalent Bonds
Suppose two atoms both attract electrons unequally, but with not such a big difference that a transfer of electrons occurs making the bond ionic. The resultant bond is called a polar covalent bond. A good example is the hydrogen chloride molecule, HCl.

Binary Covalent Compounds
Compounds between non-metals have covalent bonds. They are not ionic, and you cannot determine their formulas from ionic charges. Such compounds are named with prefixes denoting the number of each atom in the compound.

Phosphorus trichloride
PCl3

Phosphorus pentachloride
PCl5

Nitrogen dioxide
NO2

Dinitrogen pentoxide
N2O5

Sulfur hexafluoride
SF6

Ionic vs. Covalent Compounds—Different Properties
Ionic compounds have an ionic lattice—strong ionic bonds throughout the entire solid in a threedimensional array. Most ionic compounds have high melting points and boiling points. Covalent compounds have strong bonds within the molecule (covalent) but much weaker bonds between molecules (intermolecular forces). We will be studying intermolecular forces in a future chapter. Covalent compounds typically have much lower melting points and boiling points than ionic compounds. NaCl is a solid at room temperature, existing as an ionic lattice. Melting point: 801 oC. H2O is a liquid at room temperature. As a solid (ice), there are much weaker intermolecular forces than the ionic bonds in NaCl. Melting point of H2O (ice): 0 oC C5H12 (pentane) is an organic compound with covalent bonds. It is a liquid at room temperature, and boils when warmed slightly. Melting point of pentane: -130 oC Boiling point of pentane: 36 oC.

## Chem 161 Equations

% relative error
(|M-A|/A)100%

Average Deviation
sum|M-X|/#

R_f of Spot
(D spot traveled from origin)/(D solvent from origin)

Density
mass/volume

Ideal Gas Law
PV=nRT

Heat Required to Raise Temp
q=(C)(m)(Delta T)

% yield
(actual/theoretical)100%

Molarity
n/L

pH
-log[H+]

pOH
-log[OH-]

[OH-]
10^-pOH

[H+]
10^-pH

K
[P]/[R]

## Chem 161 Compounds

Covalent Boding
sharing of electrons

Ionic Bonding
transfer of electrons

Covalent Bonding occurs between…
non-metals

Ionic Bonding occurs between…
metals and non-metals

Which electrons are involved in bonding?
valence

In Ionic Bonding, the metal becomes…
a cation (+)

In Ionic Bonding, the non-metal becomes…
an anion (-)

Ions generally arrange into…
a crystal lattice as solids

Molecular Formula
shows the exact number of atoms of each element in a molecule

Empirical Formula
The simplest whole number ratio of atoms of each element present in a compound

Unit Formula
The empirical formula of an ionic compound

Ionic Compounds must be electrically…
neutral

Order for writing empirical formulas
C,H,N/O, then alphabetical

To name a monoatomic anion
add -ide to the element’s name

To name a monoatomic cation
add -ion after the name of the element

Metal with a charge of 2
-ous

Metal with a charge >2
-ic

Hydronium
H3O +

Ammonium
NH4 +

Mercury(i)
Hg2 2+

Hypochlorite
ClO –

Chlorite
ClO2 –

Chlorate
ClO3 –

Perchlorate
ClO4 –

Chlorates belong to group

## Chem 161 Polyatomic Ions

ammonium
NH4+

Acetate
C2H3O2 −

cyanide
CN-

hypochlorite
ClO-

chlorite
ClO₂-

chlorate
ClO₃-

perchlorate
ClO₄-

dihydrogen phosphate
H₂PO₄-

hydrogen carbonate or bicarbonate
HCO₃-

hydrogen sulfate or bisulfate
HSO₄-

hydroxide
OH-

permanganate
MnO₄-

nitrite
NO₂-

nitrate
NO₃-

carbonate
CO₃²-

chromate
CrO₄²-

dichromate
Cr₂O7²-

peroxide
O₂²-

hydrogen phosphate
HPO₄²-

sulfite
SO₃²-

sulfate
SO₄²-

phosphate
PO₄³-

hydrogen sulfite or bisulfite
HSO3-

## CHEM 161 EXAM 1 Answers

scientific theory
explains how/why. based on observations/experiments, makes testable predictions, changes over time based on new evidence, can be falsifiable

scientific law
predicts what happens

intensive property
independent of the amount of substance present (color, odor, maleability, density)

extensive property
depends upon amount of substance you have (mass, length, volume)

physical property
A characteristic of a pure substance that can be observed without changing it into another substance (physical state, boiling point, color, smell)

chemical property
A characteristic of a pure substance involving its chemical change (ability to react with oxygen, ability to react with water)

mixtures
material that can be separated by physical means into two or more substances (bronze, salt water)

physical change
change in the form of matter but not in its chem identity (boiling, melting, dissolving)

chemical change
equals chem reaction, change in which one or more kinds of matter are transformed into a new kind of matter or multiple new kinds of matter (rust, photosynthesis, digestion)

compound
substance composed of characteristic proportions of two or more elements chemically bonded together

heat
transfer of energy from one object to another due to a difference in temperature

sig figs
All the digits that can be known precisely in a measurement, plus a last estimated digit

the rule that the result of a calculation is known only as well as the least well-known value used in the calculation

precision
closeness of the set of values obtained from repeated measurement of the same quantity

accuracy
the closeness of a single measurement to its true value

what digits are significant
the digits that represent the value and precision of the measurement are significant. all non-zero digits, zeros between non-zero digits, trailing zeros in a number containing a decimal point, zeros following digits following a decimal point

exact numbers
infinite number of sig figs. usually counted, unmeasured, or pre-defined numbers

sig figs in multiplication and division
Round the answer to the same number of sig figs as the measurement with the least number of sig figs used in the calc

sig figs in addition and subtraction
the number of digits is limited by the number in the calc with the smallest number of decimal places

mega
10^6 (M)

kilo
10^3 (k)

deci
10^-1 (d)

centi
10^-2 (c)

milli
10^-3 (m)

micro
10^-6 (u)

nano
10^-9 (n)

pico
10^-12 (p)

Angstrom
10^-10 (m)

molecule
two or more atoms held together by covalent bonds

density
Density = mass/volume. no matter how big/small the volume, the density of the same substance remains the same

sig figs
All the digits that can be known precisely in a measurement, plus a last estimated digit. – non-zero numbers are always sig – zeros between non-zero numbers are always sig – leading zeros are never sig – zeros at the end of a number only if the number has a decimal

the rule that the result of a calculation is known only as well as the least well-known value used in the calculation

Dalton’s Atomic Theory
1) elements are composed of atoms. 2) atoms of same element are identical, but differ from other elements. 3) elements can mix together 4) atoms only change when mixed with other elements

Electron
A subatomic particle that has a negative charge. joseph john Thompson, 1898, cathode ray tube (discovering the electron, glowing rays attracted to the positive electrode)

JJ Thomson’s atomic model
– atoms are made of smaller particles – atoms are electrically neutral – electrons are balanced by a positive charge – plum pudding model – used cathode ray tube to discover this

mililikans oil drop experiment
– robert millikan – determined charge and mass of an electron – able to find the charge of electron and then its mass – found that particles can gain and lose electrons

Rutherford’s Gold Foil Experiment
– nuclear model – ernest rutherford – was testing the plum pudding model by investigating alpha particles – 1 in 8000 particles were deflected by more than 90 degrees which was at odds with the plum pudding model – atom is mostly empty space – positive charge is centrally located in a dense nucleus (most of atoms mass) – nucleus is 100,000 times smaller than the atom – electrons are located at a location far away from the nucleus

nuclear model redefined
rutherdord knew the nucleus was more massive than the mass from protons alone – must be another particle. in 1932, james chadwick discovered the subatomic particle known as the neutron – no charge

problem with the planetary model
electrons orbiting the nucleus should radiate energy, slow down, be pulled into the nucleus and collapse the atom but this is not the case because atoms are stable and don’t collapse

atomic number
the number of protons in the nucleus of an atom

mass number
the sum of the number of neutrons and protons in an atomic nucleus

isotope
Atoms of the same element that have different numbers of neutrons. atomic number determines the element

average atomic mass
weighted average calculated by multiplying relative abundances by their atomic masses and summing the product. is the sum of the relative abundance x the atomic mass

nuclear stability
an atom that is electrically neutral has an equal number of protons and electrons. as the number of protons increases, more neutrons are needed to get a stable nucleus

Metals
Elements that are good conductors of electric current and heat. shiny, malleable, ductile solids

nonmetals
Elements that are poor conductors of heat and electric current

metalloids (semimetals)
elements that have physical properties of metals anad chemical properties of nonmetals

average atomic mass
the weighted average of the atomic masses of the naturally occurring isotopes of an element. calculate by multiplying the natural abundance of each isotope by its mass in amu and then summing the products

natural abundance
the relative percentage of a particular isotope in a naturally occurring sample with respect to other isotopes of the same element

molecular mass
mass in amu of one molecule of a molecular compound

molar mass
the mass of one mole of a pure substance

groups on periodic table
a vertical row of elements in the periodic table

periods on periodic table
A horizontal row of elements in the periodic table

group 1
alkali metals

group 2
alkaline earth metals

group 8
noble gases

group 7
Halogens

Group 3-12
transition metals

Hydrogen (H) ion (group 1)
H+

lithium (Li) ion (group 1)
Li+

sodium (Na) ion (group 1)
Na+

potassium (K) ion (group 1)
K+

rubidium (Rb) ion (group 1)
Rb+

caesium (Cs) ion (group 1)
Cs+

beryllium (Be) ion (group 2)
Be2+

magnesium (Mg) ion (group 2)
Mg2+

calcium (Ca) ion (group 2)
Ca2+

Strontium (Sr) ion (group 2)
Sr2+

Barium (Ba) ion (group 2)
Ba2+

aluminum (Al) ion (group 3)
Al3+

gallium (Ga) ion (group 3)
Ga3+

indium (In) ion (group 3)
In3+

nitrogen (N) ion (group 5)
N3-

oxygen (O) ion (group 6)
O2-

Sulfur (S) ion (group 6)
S2-

selenium (Se) ion (group 6)
Se2-

Tellurium (Te) ion (group 6)
Te2-

flourine (F) ion (group 7)
F-

chlorine (Cl) ion (group 7)
Cl-

Bromine (Br) ion (group 7)
Br-

iodine (I) ion (group 7)
I-

1 mole of an element has how many atoms
6.022×10^23 atoms

1 mole of a compound has how many molecules
6.022×10^23 molecules

molarity
M = moles of solute/liters of solution

solution
A homogeneous mixture of two or more substances. the substance in the greatest amount is the solvent

aqueous solution
a solution in which water is the solvent

concentrations
quantity of solute dissolved in a specific quantity of solvent or total solution (amount of solute/amount of solution)

Dilutions
M1V1=M2V2 A new solution of less concentration created by adding more solvent to the original solution.

solute
A substance that is dissolved in a solution.

solvent
the substance in which the solute dissolves

isotopes
Atoms with the same number of protons but different numbers of neutrons. have different atomic masses

relative abundance of an isotope
the percentage of atoms with a specific atomic mass found in a naturally occurring sample of an element

mass number
the sum of the number of neutrons and protons in an atomic nucleus
isotopic notation
notation that shows the chemical symbol, atomic number, and mass number for an isotope of an element
element notation on periodic table

what is smaller than an atom
atom then chemical bonds then molecules then IMFs
Rutherford’s model of the atom
proposed that most of the mass of the atom was concentrated at the atom’s center (contains a dense nucleus). has an orbiting, fixed electron. is mostly empty space. diameter of atom, is 10^-10m and the nucleus diameter is 10^-15m

problems with rutherfords model
does not obey classic laws of physics. atoms are stable, the electrons do not spiral into the nucleus.

wave theory
light is a wave. Robert Hooke and Christian Huygens

Particle Theory
Light is made up of tiny particles. Isaac Newton and Pierre la Place

evidence for light being a wave
proved by experiment by Thomas Young 100yrs later. light interferes with itself like a water wave does

light waves refract through a prism
white light is composed of a continuous spectrum of visible electromagnetic radiation

wavelength
distance between wave crests

frequency
cycles per second

electromagnetic wave theory
experiments by Michael Faraday. theory by James Clerk Maxwell. electromagnetic waves have a variety of frequencies and wavelengths, but all travel at the speed of light.

electromagnetic spectrum experiments
photoelectric effect blackbody radiation emission spectra of atoms

light emitted by black body when heated up (stove) blackbody radiators absorb all light (look black), are at the same temp as surroundings, emit light when heated, can measure intensity and wavelength of emitted light results of experiment show that as temp increases, light gets brighter peak wavelength gets shorter as temperature increases predictions from wave theory dont match experiment

planck’s quantum theory
measured blackbody radiation. did not follow wave theory. proposed a new theory of light – light is absorbed and given off not ad a continuous wave but in little packets of light energy. quanta of energy planck’s quantum theory of light was a turning point in physics

photelectric effect
only light frequencies above a certain threshold cause electrons to be ejected from the metal surface. conflicts with wave theory because if you keep increasing the intensity, you will eventually…

einsteins photons (photoelectric effect)
proposed that light consists of particles called photons. as Planck proposed, Einstein’s photons have a certain quanta of energy (based on frequency). his model of light explained the photoelectric effect experiments

energy of photons (photoelectric effect)
at a specific frequency (or wavelength) a photon possess a specific quantity of energy (E) in J/photon Ephoton = h(v)

## CHEM 161 EXAM 2 Answers

actual yield
the experimental quantity, in grams, of product obtained in a reaction.

the number of particles (6.022 x 10^23) of a substance in 1 mol of that substance. By definition, Avogadro’s number is equal to the number of carbon-12 atoms in exactly 12.00 g of carbon-12.

balanced
appropriate coefficients have been added such that the number of atoms of each element are the same in both reactants and products.

chemical equation
a precise quantitative description of a reaction.

coefficients
the number placed in front of the substance in a chemical equation that reflect the specific numbers of units of those substances required to balance the equation.

formula mass
the total mass of all the atoms present in the formula of an ionic compound, in atomic mass units (amu), or the mass of one mole of formula units in grams per mole.

limiting reagant
the reactant that is consumed first, causing the reaction to cease despite the fact that the other reactants remain “in excess.” Also known as the limiting reactant.

mass percent
the percent of a component by mass. mass % = (total mass component/total mass whole sub.) x 100%

molar mass
the total mass of all the atoms present in the formula of a molecule, in atomic mass units (amu) or in grams per mole.

mole
the quantity represented by 6.022 x 10^23 particles.

molecular mass
the mass of one molecule, expressed in atomic mass units (amu), or the mass of one mole of molecules in grams per mole.

percentage yield
the actual yield of a reaction divided by the theoretical yield and then multiplied by 100%.

phase
a part of matter that is chemically and physically homogeneous.

products
the substances located on the right-hand side of a chemical equation.

reactants
the substances located in the left-hand side of a chemical equation.

stoichiometric ratios
the mole ratios relating how compounds react and form products.

stoichiometry
the study and use of quantitative relationships in chemical processes.

theoretical yield
the maximum amount of any chemical, in grams, that could be produced in a chemical reaction. This value can be calculated from the equation for the reaction.

acid
a compound that produces hydrogen ions (H+) when dissolved in water.

acid-base reaction
the reaction between an acid and a base The products are water and an ionic compound.

aqueous
water-based; also implies that a dissolved substance has a sphere of hydration.

base
a compound that produces hydroxide ions (OH-) when dissolved in water.

complete ionic equation
a chemical equation that indicates all of the ions present in a reaction as individual entities.

concentration
an intensive property of a solution that describes the amount of a solute dissolved per volume of solution or solvent. The typical concentration units include molarity, ppm, and w/w.

diprotic
can produce 2 mol of H+ when it dissolves.

electrolyte
a compound that produces ions when dissolved in water.

end point
in a titration, the volume of the added reactant that causes a visual change in color of the indicator.

equivalence point
in a titration, the point at which all reactants have just been completely consumed.

half-raction
an incomplete equation that describes the oxidation or reduction portion of a redox reaction.

hydration sphere
the shell of water molecules surrounding a dissolved molecule, ion, or other compound. This shell arises because of the force of attraction between the water molecules and the solute.

insoluble
not capable of dissolving in a solvent to an appreciable extent.

molar (M)
the “shorthand” method of describing molarity, as in “that is a 3 molar HCl solution.”

molarity (M)
a specific concentration term that reflects the moles of solute dissolved per liter of total solution.

molecular equation
a chemical equation that shows complete molecules and compounds.

monoprotic
can produce 1 mol of H+ when it dissociates.

net ionic equation
a complete ionic equation written without the spectator ions.

nonelectrolyte
a compound that doesn’t dissociate into ions when it dissolves.

oxidation
the process of losing electrons. Such a substance is sad to be oxidized.

oxidation number
a “bookkeeping” number that reflects the charge on an ion.

oxidation-reduction reaction
reactions that involve the transfer of electrons from one species to another. Also known as redox reactions.

oxidation state
oxidation numer

oxidized
the species that has lost electrons in a redox reaction.

parts per billion (ppb)
one gram of solute per billion grams of solution.

parts per million (ppm)
one gram of solute per million grams of solution.

parts per trillion (ppt)
one gram of solute per trillion grams of solution.

precipitation reaction
a reaction involving the formation of a sold that isn’t soluble in the reaction solvent.

precipitation
any solid material that forms within a solution; the action describing the formation of a solid.

redox reactions
oxidation-reduction reactions

reduced
the species that has gained electrons in a redox reaction.

reduction
the process of gaining electrons. Such a substance is said to be reduced.

soluble
the ability of a substance to dissolve within a solution.

solvent
a compound that typically makes up the majority of a homogeneous mixture of molecules, ions, or atoms; dissolves the solute.

spectator ions
ions that do not participate in a reaction.

strong acid
an acid that completely dissociates in solution.

strong base
a base that completely dissociates in solution.

strong electrolyte
any compound that completely dissociates in solution.

standard solution
a solution with a well-defined and known concentration of solute.

titration
the process of adding one reactant to an unknown amount of another until the reaction is complete; used to determine the concentration of an unknown solute.

triprotic
can produce 3 mol of H+ when it dissolves.

weak acid
an acid that partially dissociates in solution.

weak base
a base that partially dissociates in solution.

weak electrolyte
any substance that only partially dissociates in solution.

activation energy
the energy that is required to indicate a chemical reaction.

biomass conversion
the release of energy from the chemical conversion – often simply the burning – of plants and trees.

bomb calorimeter
apparatus in which a chemical reaction occurs in a closed container, allowing the energy released or absorbed to be measured.

Calorie
(C) Unit of energy equal to 1000 calories (1 kilocalorie) and to 4184 joules.

calorie
(c) Unit of energy equal to 4.184 joules.

calorimetry
the study of the transfer of heat in a process.

chemical energy
energy that is stored in a substance as a result of the motions and positions of its atomic nuclei and their reactions.

constant-volume calorimetry
a form of calorimetry in which the reacting system is sealed within a chamber of fixed volume, and the only way the system can release or gain energy is by the exchange of heat with the surroundings.

energy that propagates through space. Examples include visible light, X-rays, and radiowaves.

electroweak force
the fundamental force responsible for the attraction between objects carrying opposite electric charges, for the repulsion between objects carrying the same electric charge, and for some transformations within subatomic particles. It is also responsible for the phenomena of magnetism and light.

endothermic reaction
a reaction that absorbs energy from the surroundings.

enthalpy
a thermodynamic quantity symbolized by H and defined as H = U + PV.

exothermic reaction
a reaction that releases energy into the surroundings

first law of thermodynamics
the total change in the closed system’s energy in a chemical process is equal to the heat flow (q) into the system and the work one (w) on the system: difference in U = q + w. Energy is neither created nor destroyed; it is only transferred from place to place an converted from one form into another: diff U systems + diff U surroundings = 0.

geothermal systems
power-generating systems that involve drilling deep into the earth and exploiting the flow of heat from the interior of the earth out toward the surface.

gravitational force
the fundamental force that causes all objects with mass to be attracted to one another.

heat (q)
the energy that is exchanged between a system and its surroundings because of a difference in temperature between the two.

heat capacity
the amount of heat needed to raise the temperature of any particular amount of a substance by 1 degree celsius.

heat of reaction
the energy as heat released or absorbed during the course of a reaction.

Hess’s law
thermodynamic law stating that the enthalpy change of a chemical reaction is independent of the chemical path or mechanism involved in the reaction. This enables us to determine the enthalpy changes of reactions that might be very difficult to actually perform.

hydroelectric systems
power-generating systems in which the power of falling water is used to generate electricity.

internal energy (U)
the energy of a system defined as the sum of the kinetic and potential energies. Absolute internal energy is difficult to measure.

joule (J)
the SI unit of energy. In terms of base units, 1 Joule = 1 kg x m^2 x s^-2

kilocalorie
unit of energy equal to 1000 calories and to 1 Calorie.

kinetic energy
the energy things possess as a result of their motion.

law of conservation of energy
law stating that energy is neither created nor destroyed but is only transferred from place to place or converted from one form into another.

molar heat capacity
the heat capacity of one mole of a substance.

photovoltaic systems (PVs)
fabricated systems that convert the energy of sunlight into electricity. Also known as solar cells.

potential energy
the energy things possess as a result of their positions, such as their position in a gravitational or electromagnetic field.

reference form
the most stable form of the element at standard conditions.

renewable sources of energy
energy sources that can be rapidly replaced by natural processes.

solar cells
fabricated systems that convert the energy of sunlight into electricity. Also known as photovoltaic systems or PVs.

solar thermal systems
fabricated systems that concert the energy of sunlight into heat.

specific heat
the amount of heat needed to raise the temperature of one gram of a substance by one degree Celsius (or one kelvin) when the pressure is constant. Also known as the substance’s specific heat capacity.

specific heat capacity (c)
the amount of heat needed to raise the temperature of one gram of a substance by one degree Celsius (or one kelvin) when the pressure is constance. Also known as the substance’s specific heat.

standard enthalpy of combustion
the enthalpy change when one mole of a substance in its standard state is completely burned in oxygen gas. Also known as the substance’s standard head of combustion.

standard enthalpy of formation
the enthalpy change for the formation of one mole of a substance in its standard state from its elements in their reference form. Also known as the standard heat of formation.

standard enthalpy of reaction
enthalpy change of a reaction in which all of the reactants and products are in their standard states.

standard state
the state of a chemical under a set of standard conditions, usually at 1 atmosphere of pressure and a concentration of exactly 1 molar for any substances in solution. Standard states are often reported at 25 degrees celsius.

state function
a property of a system that depends only on its present state, not on the path by which it reached that state.

strong nuclear force
the force that holds protons and neutrons together in atomic nuclei.

surroundings
everything in contact with a chemical system. Energy may flow between the surroundings and systems. Strictly speaking, the surroundings of a chemical system comprise the entire universe except the system.

system
any set of chemicals whose energy change we are interested in.

thermochemistry
the study of energy exchange.

universe
the space consisting of both the system and the surroundings.

wind power
the use of the energy of the wind to generate electricity.

work (w)
force acting on an object over a distance.

## Chem 161 Final Exam Answs

Beta particle
emission that is a high energy electron. the electron comes from a neutron so when it is emitted, the atomic number goes up by one, but the atomic mass remains the same

alpha particle
emission that is basically a helium nucleus (2+ charge), so when it is emitted the atomic number goes down by 2 and the atomic mass goes down by 4

Alkali metals
first column

Alkaline earth metals
second column

halogens
group before the noble gases (group 17)

Naming ionic compunds
keep the name of the cation (first element) then add -ide to the anion

acetate
CH3COO(-)

Ammonium
NH4(+)

Azide
N3(-)

Carbonate
CO3(2-)

Chlorate
ClO3(-)

Chromate
CrO4(2-)

Cyanide
CN(-)

Dichromate
Cr2O7(2-)

Dihydrogen phosphate
H2PO4(-)

Hydrogen carbonate/bicorbonate
HCO3(-)

hydrogen phosphate
HPO4(2-)

Hydrogen sulfite/bisulfite
HSO3(-)

Nitrate
NO3(-)

Nitride
N(3-)

Nitrite
NO2(-)

Perchlorate
ClO4(-)

Permanganate
MnO4(-)

Peroxide
O2(2-)

Phosphate
PO4(3-)

Sulfate
SO4(2-)

Sulfite
SO3(2-)

Thiocyanate
SCN(-)

Hypochlorous acid
HClO

chlorous acid
HClO2

chloric acid
HClO3

perchloric acid
HClO4

Spectrophotometry: Beer’s Law and molar absorptivity
A = εbc A (absorbance) is a measure of color intensity ε is a constant f molar absorptivity b is the path length the light travels through the solution c is the concentration

Strong electrolyte
Compound that dissociates completely into ions in solution, enhancing the conductivity of the solution

Acid is a…
Proton donor (HCl)

Base is a…
Proton acceptor (NaOH)

Hydrolysis
reaction of water with another material. Hydrolysis of nonmetal oxides produces acid

All compounds containing _______ are soluble
Group 1 ions, NH4(+), NO3(-), and CH3OO(-)

All group ____ ions are soluble except those containing ______
All group 17 are soluble except those containing Ag, Cu, Hg2, and Pb (a.k.a. the heavy metals)

SO4(2-) is always soluble unless is contains ______
Ba, Ca, Hg2, Pb, and Sr

Insoluble Compounds
All hydroxides, sulfides, carbonates, phosphates, and flourides except all of those with group 1 cations and Ca, Sr, Ba, and NH4(+)

Oxidizing agent
The substance in a redox reaction that contains the element being reduced

Reducing agent
The substance in redox reaction that contains the element being oxidized

Steps for balancing redox in acidic solution
1) Create the half-reactions 2) Balance the elements except those containing oxygen or hydrogen 3) Balance the oxygens with water 4) Balance the water with H+ on the other side of the equation 5) Balance the charges with electrons 6) Balance the electrons in each equation by multiplying by some whole number 7) Add the two half-reactions

Molar heat of fusion (ΔHfus)
q = nΔHfus

Molar heat of vaporization (ΔHvap)
q = nΔHvap

Units of heat capacity
kJ/molC

Units of ΔHrxn
kJ/mol

Determining ΔHrxn using calorimetry
Figure out the moles of water in the reaction (whether that be using density given volume or molar mass given mass), then calculate the q of the water given the heat capacity of water, the change in temperature, and the moles. Then that number equals the amount of energy given off by the reaction (negative value). Plug these in and solve for whatever unknown

ΔHrxn equation
ΔHrxn = Σ n.productsΔHf.products – Σ nreactantsΔHf.reactants

Steps for determining ΔHrxn using ΔHformation
1) Write the formation equation for every substance (reactants and products) that isn’t an element 2) Flip each equation so that the reactants are on the reactant side and the products on the product side (multiply each ΔHf by -1) 3) Multiply each equation and ΔHf by the stoichiometric coefficient in the balanced equation that you’re producing. 4) Add em up

heating curve of water
the sloped lines represent one state of water. the flat lines represent when it is in both states in a phase transition (the temperature does not change even though heat is being added)

Hess’s Law
States that the standard enthalpy of reaction ΔH.rxn for a reaction that is the sum of two or more reactions is equal to the sum of the ΔH.rxn values of the constituent reactions

Relationship between r & l of the simple cubic unit cell
r = l/2 where l is just the length of one side and is 2 radii long

Relationship between r & l of the body-centered cubic unit cell
r = l sqrt(3)/4 where lsqrt(3) is the 3d diagonal, which is 4 radii long

Relationship between r & l of the face-centered cubic unit cell
r = l sqrt(2)/4 where l sqrt(2) is the diagonal of one face which is 4 radii long

Raoult’s Law
P.solution = X.solvent P.solvent the vapor pressure of the solution is equal to the mole fraction of moles solvent/moles solution times the vapor pressure of the solvent i.e. P.seawater = (moles water/moles seawater)(vapor pressure of the water)

Henry’s Law
Solubility of gas in liquid k C.gas = k.h P.gas The concentration of a sparingly soluble gas in liquid is proportional to the partial pressure of the gas in the environment surrounding the liquid

If you are given pressure and need to find concentration use ______
Henry’s law and multiply by the Henry’s constant

ΔH.solution when dissolving an ionic compound in solution
ΔH.solution = ΔH.hydration – U

Born-Haber Cycle
1) Sublimation: turning the elemental solid into gas ΔH.sub 2) Break all covalent bonds in the gas phase to create 1 mole of the element (i.e. 1/2Cl2(g) -> 1 Cl(g)) using ΔH.BE (where in the case of 1/2Cl2 it would be 1/2ΔH.BE because its have a mole of Cl2 bonds) 3) – Ionization of the element -> cation in the gas phase using IE1 or IE2 respectively – Ionization of the element -> anion using EA 4) Formation of 1 mole of the ionic compound where ΔH.lattice = U which is the lattice energy

Electron configurations (f rows, d rows, and the order of the orbitals)
14 electrons in f that start with 4 then 5, 10 in d that starts with 3 to 6, and 2 s normally. You order it in increasing n

Strength of IMFs (from strongest to weakest)
ion-ion, ion-dipole, hydrogen bonding, dipole-dipole, dipole-induced dipole , dispersion

Energy of a reaction (E not ΔH!)
q (J) = nCΔT

How do you figure out the total energy given off in a reaction?
ΔH x moles of reactant = q Then if you want to find the change in temperature, you plug that into q = nCΔT

General phase diagram
Most left solid, bottom gas, top middle liquid, top right corner supercritical fluid

Molality
moles of solute/kg of solvent

Boiling/freezing point elevation/depression
ΔT = ikm

Osmotic pressure
Π = iMRT

Cubic closest packed unit cell type 2 Types of cubic packing unit cell types
ccp –> fcc cp –> bcc/sc

Calculate density of an element given unit cell info and atomic mass
First draw a picture Find out how many atoms there are per unit cell, then find out how many moles there are given that many atoms (really small number), then use atomic mass to convert to grams per unit cell. Then use the relationship between r and l to find l and cube it to find the volume and convert to cm^3 using the conversion. Then divide mass by volume.

## Chem & 161 Course Overview

This course is required for most students intending to go into sciences or engineering and pre-professional health students. This intro-level chemistry course should be a breeze if you have already taken an introductory chem class in either high school or college.

### Learning Goals

• Understand and apply basic principles and concepts in the physical or biological
sciences;
• Explain and be able to assess the relationship among assumptions, methods, evidence, arguments, and theory in scientific analysis;
• Department Learning Goals Met by this Course.

By the end of this course, students will be able to draw upon:

• relevant scientific models;
• representations at the macroscopic, submicroscopic (small particle), and symbolic levels—including mathematical formulae;
• qualitative and quantitative reasoning skills.  