Honors Chemistry Units 1A-3 Review

Honors Chemistry Units 1A-3 Review

UNIT 1A

Lab equipment

Beaker: Used for heating, carrying and stirring things. NOT used for measuring.
Beaker tongs: Used to carry beakers.
Graduated Cylinder: Used for measuring the volume of liquids. Make sure to read from the bottom of the meniscus.
Erlenmeyer Flask: Used for collecting liquids and mixing solutions. Designed for swirling things.
Test Tube: Holds liquids or solids
Test Tube Holder: Holds test tubes to heat them above flames
Volumetric Flask: Making a specific volume of a solution
Buret: Used for dispensing solutions, had graduations
Glass stirring rod: Used for stirring things, helps guide the flow of a liquid.

Beaker

Beaker tongs

Graduated cylinder

Erlynmeyer flask

TEst tube

Test tube holder

Buret

Glass stirring rod

Volumetric flask

Lab equipment con't

Mortar and Pestle: Used for grinding, bowl is the mortar, pestle is the grinder
Evaporating dish: Evaporating water/liquids from a mixture
Watch Glass: Holds solids, evaporating, cover glassware, keeps liquid in evaporating dish from splatteirng
Crucible: Used for heating, lid controls airflow
Crucible Tongs: Carrying a crucible
Clay Triangle: Holding crucible for heating
Bunsen Burner: Heats glassware
Striker: Lights bunsen burner
Ring stand: Holds clamps

Mortar and Pestle

Evaporating Dish

Watch glass

Crucible

Crucible tongs

Clay triangle

Striker

Ring stand

Bunsen burner

Lab equipment Part 3

Iron ring: Supports objects for heating
Utility clamp: Holds test tubes/thermometers on ring stand
Wire Guaze: Holds glassware above flame
Pipet: Transferring solutions, adding drops of solutions
Scoopula: Scooping solids, transferring solids
Well plate: Performing small scale tests
Thermometer: Measuring temperature

Iron Ring

Utility Clamp

Wire guaze

Pipet

Scoopula

Well plate

Thermometer

Significant Figures

• Calculations can only be as precise as our least precise measurement

• Significant Figures (Sig Figs) are used to:

• Show precision of measurement

• Round answers to 3 sig figs. 

  There are unlimited significant figures when it comes to: counting (ex: 24 people in a classroom), exact, nonmeasured quantities, ex: 60 minutes = 1 hour

Multiple Choice

Is this a situation where there are unlimited significant figures?

8 rubber ducks

• No, this is not a situation with unlimited significant figures.
• Yes, this is a situation where there are unlimited significant figures.

WHy yes?

It demonstrates a situation with counting, and we established in slide 11 that when counting is involved, there are unlimited significant figures.

Multiple Choice

Is this a situation where there are unlimited significant figures?

583 ms

• No, this is not a situation with unlimited significant figures.
• Yes, this is a situation where there are unlimited significant figures.

WHy no?

We can actually apply our sig fig rules to this problem.

The number is 583, it has 3 significant figures in total because all numbers 1-9 are always significant.

Significant figure rules

1: All numbers 1-9 are significant ex: 583 ms (3 sig figs)
2: Interior zeroes (zeroes between nonzero #s) are significant ex: 909 (3 sig figs)
3: Leading zeroes before the first nonzero are never significant ex: 0.001 km (1 sig fig)
4: Trailing zeroes after the last nonzero digit are only significant when a decimal point is in the number ex: 900. kg (3 sig figs)

Multiple Choice

How many significant figures are in this number?

2.21

• 2
• 3
• 4
• 5

w hy 3?

All numbers 1-9 are significant.

Multiple Choice

How many significant figures are in this number?

4,000,000

• 2
• 1
• 3
• 5

w hy 1?

Trailing zeroes are only significant if there is a decimal point. In this situation, we simply have commas so its just 1.

Multiple Choice

How many significant figures are in this number?

0.082900

• 4
• 5
• 6
• 3

w hy 5?

The first 2 zeroes aren't significant becuase they're leading zeroes and those are NEVER significant. 829 are significant because they're numbers 1-9 and the last 2 zeroes are due to the decimal point.

Significant figure rounding

For rounding:
Find the number that represents the last significant figure you need, 5 goes up, 4 goes down

Examples:
8.000 ---> 8.00
24.63 ---> 24.6

Multiple Choice

Round this significant number:

9.196

• 9.19

• 9.10

• 9.20

• 9.2

Why 9.20?

Find the number that represents the last significant figure you need, 5 goes up, 4 goes down

In this case it was 6, 5 or more goes up so 19 becomes 20.

Dimensional analysis

Dimensional Analysis is a method of converting one unit to another one.

Conversion factors:

The ratios between 2 units.
Always equivalent to 1.

Ex: 12 in = 1 ft, so: 12 in/1ft or 1 ft/12 in

Always make sure to write out units!

Example problem: How many quarters are in 3 dollars?

1. Do quarters/dollars so we can cancel out the 3

2. 3 dollars/1 quarter = 4 quarters/1 dollar

3. Answer: 12 quarters

Example problem 2: Your friend weighs 71.8 kg. Knowing that there are 2.2 lbs in 1 kg, how many lbs does he weigh?

1. Set up the equation so it's 71.8 kg/1 lb = 2.2 lbs/1 kg

2. Answer: 158 lbs

   So the rule of thumb is pretty much to make sure the unit on the bottom of the 2nd equation matches the one in the first so you can cancel out the kgs.

Multiple Choice

You try: You need to measure out 38.1 g of CH4. If 1 mole of CH4 has a mass of 16.05 grams, how many moles of CH4 do you measure?

• 2.37 moles
• 1.50 moles
• 3.00 moles
• 0.95 moles

Why 2.37?

The setup to your problem should be like this:
38.1 grams/1 = 1 mole/16.05 grams
You get 38.1/16.05 = 2.37 moles

Multiple Choice

You try:  If you run a 5.00 km race, how many feet do you travel? (1 mile = 1.6 km, 1 mile = 5280 ft)

Note you have to do more than 1 conversion!

• 16500 ft
• 20000 ft
• 18000 ft
• 15000 ft

Why 16,500?

The setup to your problem should be like this:
5.00 km/1 = 1 mile/1.6 km
This gets you 312.5 miles/1 = 5280 ft/mile

Multiple Choice

How many feet long is a 341 cm cable? (1 in = 2.54 cm) (1 ft = 12 in)

• 40.5 ft

• 341 ft

• 9.5 ft

• 11.2 ft

Why 11.2?

The setup to your problem should be like this:
341 cm/1 = 1 inch/2.54 cm = 134.3 in/1 = 1 ft/12 in
This gives you 11.2 ft

SI Units

• SI (AKA Metric System)

• Used by the rest of the world (incl. science)

• Why?

• It's universal (used everywhere but US)

• Units are converted based on factors of 10 

What are the SI units for:

•  Length/distance? Meters (m)

• Time? Seconds (s)

• Temperature? Kelvin (K), the absolute scale for Celsius (degrees C)

• Mass? Kilograms (kg), grams are base unit

What are the SI units for:

• Amount of substance?

• Mole (mol)

• Electric current?

• Ampere (A)

• Luminous Intensity?

• Candela (cd)

Prefixes

• Identify the magnitude of the measurement

• Examples: kg, cm, mL

SI Unit conversion examples

We must always compare the unit to the base unit
Ex: 1 centimeter is 10^-2 m this means that 10² cm = 1 meter
Ex 2: 1 kilometer is 10³ m this means that 10³ km = 1 meter

Use this relationship to make the conversion factor
So 10² cm = 10^-3 km

Basically change the sign of the magnitude!

Example problem

We want to know 100 mm is what value in cm?
The setup is as follows:
100 mm/1 = 10² cm (change from - to +)/10³ mm (change - to + and keep on bottom to cancel out) this gives us 10 cm in total, use calculator to do the math.

Use this chart to help!

Multiple Choice

500 mL = ? hL

• 5.5 hL

• 500 hL

• 0.00500 hL

• 880 hL

Why 0.00500?

The setup to your problem should be like this:
500 mL/1 = 10^-2 hL (hecto is 10² but we change sign)/10³ mL (change from - to +)
Use calculator to do the math and you get 0.00500 hL

• How reproducible data is

• How close sets of data are to each other

• Requires more than 1 trial

Precision

• How close data is to an accepted value

• Evaluated with percent error

Accuracy

Accuracy vs Precision

• Errors you make in measurements or flaws in equipment

• Same effect (high/low) each time

• Avoid by identifying/fixing sources

• Describe these in error analysis of lab

• Ex: substances not fully dried, drops on side of cylinders

Systematic

• Uncontrollable fluctuations in data.

• Expected.

• Due to limitations of equipment, conditions of lab, rounding, etc.

• Equal chance of being high or low

• Can be limited by averaging measurements

• Ex: temp changes in lab, placement on balance

Random error

Types of error

​No errors: Have both accuracy + precision
Random errors: Have accuracy but not precision
Systmeatic: Have precision but not accuracy

​

Random vs systematic error

Value you get in a trial(s)

Experimental value

Value you are expecting to get
AKA: Book value, accepted value

Theoretical value

|Experimental Value - Theoretical Value|

Absolute error

Types of error

• %Error = |Experimental - Theoretical | x 100

                              Theoretical

• Absolute value, so always positive

• Accuracy means < 5% error

• Round values to 2 Decimal Places

Percent error

Example problem: In 3 straight trails, Mr. White finds the density of a silver ring to be 9.80 g/mL. When he looks up the density of silver in his chemistry book, he finds that it is 10.5 g/mL.

What is his percent error?

%Error = | 9.80 - 10.5 | x 100 , %Error = 6.67 %

                          10.5

Is his data accurate, precise, both or neither?

Precise, but not accurate.

Percent error example problem

Amount of matter in an object/sample
Never changes regardless of location/conditions

Mass (m)

Space taken up by matter
Generally increases with temp. increases
For regular solids: length x width x height (lwh)
For irregular solids: use water displacement

Volume (v)

Amount of matter in a given space
Intensive, physical property
To calculate:
D = m/V
Common units: g/mL, g/cm³, kg/m³
Note: 1 mL = 1 cm³

Density (d or p)

Density

Density

• Generally decreases as temp. increases

• Mass remains the same

• Most substances increase in volume as temperature increases

• Water is a notable exception

• Denser objects sink, less dense float (ex a bolt sinks to the bottom but a soda cap floats to the top)

• Ex: Density column

Density problem example

Problem ex: If a 45.2 g block has sides of 7.00 cm, 4.00 cm, and 1.33 cm. What is it's density?

 

D = m/V = 45.2 g/37.24 cm^3 = 1.21 g/cm^3

 

V = lwh = (7.00)(4.00)(1.33)

V = 37.24 cm^3

Density problem example 2

Problem ex 2: An empty container weighs 121.3 g. Filled with CC14 (D = 1.53 g/cm^3), the full container has a mass of 283.2 g. What is the volume of the container?

 

D = m/v                V = m/D because D and v are interchangeable.  V = 161.9 g/1.53 g/cm^3 = 106 cm^3

We got 161.9 by doing 283.2 - 121.3

Properties of matter

• All pure substances have characteristic properties

• Properties are used to distinguish between substances

• Properties are also used to separate substances

Physical property

• Characteristics observed/measured without changing a substance's composition

• Describe the substance itself

• It's red, it's a liquid, it takes up space from the half to the bottom of the beaker, pretty transparent, maybe 120 milliliters

Physical properties include:

• State of matter

• Color

• Mass, shape, length, volume

• Density

• Mallebility ---> ability to bend/hammer into thin sheet

• Ductility ---> ability to make wire

• Magnetism

• Melting/boiling point, etc

Chemical property

• Indicates how a substance will react with another

• Can be a failure to react

• Determined by changing substance's identity

• Ex: Iron rusting, silver tarnishing

Chemical properties include:

• Reactivity w/a substance

• Combustibility/Flamability ---> burn in air

• Toxicity

• Oxidation

• Decompistion

• Flammable and combustible liquids are similar but have differences

Multiple Choice

Density problem example: A 7.65 g chunk of metal is added to a graduated cylinder containing 21.0 mL of water. If the volume settles at 23.7 mL, what is the cube's density?

D = m/v

• 0.36 g/mL

• 2.83 g/mL

• 0.05 g/mL

• 0.32 g/mL

Why 2.83?

D = m/v   D = 7.65 g/2.7 mL    D = 2.83 g/mL

Pure substances

• Cannot be separated by physical means

• Every sample has the same characteristics

• Characteristics can identify substance

• AKA Chemical

Pure substances

• Elements

• Made up of ONE type of Atom

• Smallest unit that maintains the chemical identity of an element

• Smallest unit of matter with unique properties

• Listed on Periodic Table

• Ex: Carbon (C ), Nitrogen (N), Calcium (Ca), > 100 others

Pure substances

• Compounds

• Made up of two or more types of atoms chemically bonded

• Properties are different than bonded elements

• Can be broken down into elements/smaller compounds

• Must be chemical separation, not physical separation

• Ex: water (H₂0), sugar (C₁₂H₂₂₀₁₁), salt (NaCl), etc

• Ex: all green dots would be an element because it's the same, orange and yellow dots would be a compound because they're all different.

Mixture

• Blend of two or more kinds of matter (elements and/or compounds)

• Each substance retains its own identity and properties

• Anything that has mass and volume is matter

Homogeneous Mixture

• AKA Solutions

• Have uniform composition

• Same all the way through

• Particles are too small to be seen

• Examples: saltwater, tea

Alloys

• Solid solutions that contain at least 1 metal

• Blended together so they have more desirable properties

• Ex: Stainless Steel (iron/chromium/zinc) Bronze (tin/copper)

Heterogeneous mixture

• Do not have uniform composition

• You can see the particles in them

• Examples:

• Italian Dressing (oil, water, vinegar, veggies)

• Soil (dirt, rocks, worms, etc.)

Suspensions

• Heterogeneous mixture with largest particle size

• Solid particles eventually settle out of solution

• Ex: Muddy water, Mixtures of two solids, Italian Dressing

Colloid

• Heterogeneous mixture

• Won't separate upon standing

• Can't be separated by filtration

• Shows Tyndall Effect

• Scattering of light

• Ex: paint, dairy, Jello

Physical change

• Change in a substance that does not alter the substance's chemical identity (formula)

• Ex: mixing, freezing, heating, melting, volume adjustment, bending/folding it, grinding, cutting, boiling

• All changes of state are physical changes!

Changes of state

• Sublimation

• Solid becomes a gas

• Ex: dry ice, air fresheners, freeze drying

• Deposition

• Gas becomes a solid

• Ex: Iodine, frost

Chemical change

• One or more substances are converted into different substance(s)

• Always results in new substance(s) with different:

• Chemical formula(s)

• Properties

• Many (NOT ALL) are irreversible

Signs of a chemical change

• A gas is released (evolved)

• May be visible or detected as an odor

• Color change

• Only chemical if resulting solution is different color from ALL of its components

• A Precipitate is formed

• Solid falls out of solution

• Make heterogeneous from homogeneous

Signs of a chemical change pt 2

• Temperature Change

• Can also be light

• Endothermic

• Energy is absorbed by system

• Feels cold

• Ex: Instant ice pack

• Exothermic

• Energy released as heat or light

UNIT 1B

Classifying elements ---> metals

• Left side of table

• Most common class of elements

• Generally:

• Solid (except Hg (mercury))

• Conductive of heat and electricity

• Malleable (can be bent into different shape)

• Ductile (can be wires)

• Lustrous (shiny)

Classifying elements ---> Alkali metals

• Group 1 elements (Not H) (Li all the way to Fr)

• Generally dull, soft, and reactive

  • Never free elements in nature

  • React with air/water

Classifying elements ---> Alkaline earth metals

• Group 2 elements

• Harder, denser, stronger, less reactive than alkali metals

  • Can occur in nature, but are usually in compounds

Classifying elements ---> Transition metals

• Transition Metals

• Center block on Table (Groups 3-12)

Classifying elements ---> inner Transition metals

• Cerium (Ce-Lu) Lutetium

• Lanthanide Series (Rare Earth Metals)

  • Top row of f-block

• Actinide Series

• Thorium (Th-Lr) Lawrencium

• Bottom row of f-block

Classifying elements ---> Post transition metals

Post Transition Metals include: Al, Ga, In, Sn, Tl, Pb, Bi, Nh, Fl, Mc, Lv

Classifying elements ---> Metalloids

• Border the staircase

• B, Si, Ge, As, Sb, Te, Po, At, Ts

• Properties in between metals and nonmetals:

  • Brittle

  • Lustrous

  • Semiconductors

Classifying elements ---> non metals

• Right side of the staircase

• Generally:

  • Gasses or dull, brittle solids

  • Poor conductivity

  • Poor ductility

  • Non-malleable

  • Non-lustrous

 

Classifying elements ---> Halogens

• Group 17 elements (Not At)

• Most reactive nonmetals

  • As reactive as alkali metals

  • Rarely free elements

 

Classifying elements ---> Noble gases

• Elements in group 18

• Extremely unreactive and stable

  • Almost never bond

 

Classifying elements ---> representative elements

• Groups 1,2,13-18 (1A-8A)

• All elements in the same group share common characteristics

 

Note that...

• Other groups can be named by the top-most element

• Ex: Group 15

  • Nitrogen Group

 

Use this image to help!

Multiple Choice

Using the image of the periodic table what group does Iodine (I) belong to?

• Halogens

• Alkali Metals

• Alkaline Earth Metals

• Carbon Group

Multiple Choice

Using the image of the periodic table what group does Calcium (Ca) belong to?

• Alkali Metal

• Alkaline Earth Metal

• Transition Metal

• Inner Transition Metals

Multiple Choice

Using this image of the periodic table what group does Cesium (Cs) belong to?

• Alkali Metals

• Carbon Group

• Halogens

• Alkaline Earth Metals

Multiple Choice

Using this image of the periodic table what group does Argon (belong to?

• Alkali Metals

• Inner Transition Metals

• Transition Metals

• Noble Gas

Atom

• Smallest particle of an element that retains the properties of that element

• As small as ~0.5 A (angstroms)

  • 5 x 10^-11 m

• Microscopes cannot see much inside the atom

The atom

• Name: Proton, Neutron, Electron

• Proton symbol: p+  Neutron symbol: n0  Electron symbol: e-

• Charge of a proton: +1  Neutron symbol: 0   Electron symbol: -1

• Relative mass of proton: 1 Relative mass of Neutron: 1 Relative mass of electron: 0

• Location in the atom for the proton: Nucleus  Location in the atom for a neutron: Nucleus  Location in the atom for an electron: Electron Cloud

The atom

• Atomic Number (Z)

• Number of protons in nucleus

  • All protons are alike

• Same as number of electrons in neutral atom

• Determines atom's identity

Mass number

• Sum of protons + neutrons in nucleus

  • Essentially all mass of atom

• Note: DO NOT ROUND NUMBER ON PERIODIC TABLE!

  • Will be given to you or determined from problem

Relationship between mass # and protons, neutrons, and electrons

• Protons (p+)

• Equal to Atomic Number

  • Neutrons (n0)

• Mass Number - Atomic Number

  • Electrons (e-)

• Equal to atomic number in neutral atom

Isotope

• Specific atom of an element

  • Most elements have 2+

• Isotopes of same element have:

  • Same number of protons/electrons

  • Different numbers of neutrons

  • Different mass numbers

Isotope

• To name:

  • Always include mass number

  • Put mass number after name of element

• Carbon-12, Carbon-14, etc.

Element symbols

• Symbol of the element from table

  • Note: second letter is ALWAYS lowercase

• Mass number (put it on the top left) (REQUIRED)

• Atomic number (put it on the bottom left) (OPTIONAL)

  • Optional unless "complete symbol"

Isotope examples

The isotope name is Boron-11

Our atomic number would be 5 (comes from the periodic table) Mass number would be 11 (this comes from the isotope NOT THE PERIODIC TABLE!) Number of protons would be 5 (has to match with atomic number) Number of neutrons would be 6 (11 -5 = 6) and number of electrons would be 5 (has to match with protons + atomic number)

Isotope Name:  Aluminum-27  
Atomic number: 13   Mass Number: 27  Number of Protons: 13 Number of Neutrons: 14 Number of Electrons: 13

Multiple Choice

You try! Given this Isotopes name: Ru-97 find atomic number, mass number, number of protons, number of neutrons, and number of electrons

Use the table to help

• Atomic Number: 45, Mass Number: 97, Protons: 45, Neutrons: 51, Electrons: 45

• Atomic Number: 44, Mass Number: 97, Protons: 44, Neutrons: 53, Electrons: 44

• Atomic Number: 44, Mass Number: 97, Protons: 44, Neutrons: 51, Electrons: 45

• Atomic Number: 43, Mass Number: 98, Protons: 43, Neutrons: 55, Electrons: 43

How we got that answer

The isotope name is Ru-97

Our atomic number would be 44 (comes from the periodic table), Mass number is 97 (given), Protons and Electrons would be 44 as they both have to match the atomic number, Neutrons would be 53 (97 - 44)

Multiple Choice

You try! Given this Isotopes name: U-235 find atomic number, mass number, number of protons, number of neutrons, and number of electrons

Use the table to help

• Atomic number: 235

  Mass number: 235

  Number of protons + electrons: 235

  Number of neutrons: 0

• Atomic number: 81

  Mass number: 235

  Number of Protons + Electrons: 81

  Number of neutrons: 154

• Atomic Number: 92

  Mass Number: 235

  Number of Protons + Electrons: 92

  Number of neutrons: 0

• Atomic number: 92

  Mass number: 235

  Number of protons + electrons: 92

  Number of neutrons: 143

How we got that answer

The isotope name is U-235

Our atomic number would be 92 (comes from the periodic table), Mass number is 235 (given), Protons and Electrons would be 92 as they both have to match the atomic number, Neutrons would be 143 (235 - 92)

Ions

• What if atoms aren't neutral?

• Ions

• Charged atoms

• Result from loss or gain of electrons

• Charge = p+/- e-

Anion

• Negatively charged ion

• Result from gaining electrons

• Add charge to electron total (ignore negative sign)

Ion example problem

We are given 80 (atomic mass), 34 (atomic number) and Se^2-

Number of protons: 34 (equal to atomic number)
Number of neutrons: 46 (80 - 34)
Number of electrons: 36 (34 + 2, negative means add)

Cation

• Positively charged ion

• Result from loss of electrons

• Subtract charge from electron total

Multiple Choice

You try this Ion example problem!

Identify the number of protons, neutrons and electrons given 27 (atomic mass), 13 (atomic number) and Al³⁺

• Protons: 40 Neutrons: 16 Electrons: 27

• Protons: 27 Neutrons: 13 Electrons: 27

• Protons: 13 Neutrons: 14 Electrons: 10

• Protons: 27 Neutrons: 14 Electrons: 10

How we got that answer

Protons are equal to atomic number, so 13
For neutrons we did 27 - 13 = 14
Finally for electrons we did 13 - 3 = 10 (We had a +3 so we do the opposite, subtract)

Atomic Mass

• Decimal numbers on the periodic table

• Weighted average of all isotopes of an element

• Based on abundance of each isotope in nature

• Too small to be measured in grams

  • Atomic Mass Unit (amu)

• Mass of 1 proton or 1 neutron

• 1/12th the mass of a 12^C atom

Calculating Atomic Mass

• Unless told otherwise, mass of isotope is mass number in amu

• Convert percent abundance to decimal

  • Divide by 100

• Multiply isotope mass by decimal for each isotope

• Add results

Atomic mass example problem

Neon has 3 isotopes. 20Ne has an abundance of 90.48%, 21Ne has an abundance of 0.27% and 22Ne has an abundance of 9.25%. What is the atomic mass of Neon?

20Ne: 20 amu x 0.9048% = 18.096 amu

21Ne: 21 amu x 0.0027% = 0.0567 amu

22Ne: 22 amu x 0.0925% = 2.035 amu

*All of these percentages were divided by 100%

• Final answer: 20.19 amu

• Check on periodic table to see if it's close to this result, actual atomic mass is 20.8, pretty close to 20.19

unit 2

WAVELENGTH + FREQUENCY

• Wavelength (λ) (lambda)

• Distance required for one wave cycle

• Measured in unit of distance (m, nm, etc.)

• Frequency (v)

• Number of cycles in one second

• Measured in Hertz (1 Hz = 1 cycle/sec = s^-1)

eLECTROMAGNETIC RADIATION (emr)

• AKA Light

• Vibrating particles create electric waves

  • Electric field creates magnetic field

• Energy travels as a wave through space

• Organized on Electromagnetic Spectrum

• Visible Light is only part our eyes can detect

  • Only ~.0035% of the spectrum

eLECTROMAGNETIC RADIATION (emr)

• All waves have different wavelength and frequency

• What is the relationship between wavelength and frequency? Inverse

  • Higher Frequency = Lower Wavelength

• Moves at the speed of light

  • c = Speed of Light = 3 x 10^8 m/s ---> c = wavelength x frequency for determining wavelength or frequency for any wave

1. A light has a wavelength of 5.00 x 10^-8 m.

   a. What is the frequency?

C = wavelength x frequency --> frequency = c/wavelength = 3.00 x 10⁸ m/s/5.00 x 10^-8 m/s

•  6.00 x 10¹⁵ Hz

  b. In what region of the EM spectrum is this radiation?

• UV

​

EMR problem example

1. A light has a frequency of 1.43 x 10^13 Hz

   a. What is the wavelength?

C = wavelength x frequency ---> Wavelength = c/frequency = 1.43 x 10¹³ Hz/s/3.00 x 10⁸ m/s

• 2.10 x 10^-5 m

  b. Does this type of spectrum have a shorter or longer wavelength than red light?

Shorter

​

EMR problem example 2

• AKA: Planetary Model

• Niels Bohr - 1913

• Positive nucleus at center

• Electrons orbit in energy levels

• Calculated energy released when electrons drop energy levels

• Only worked for hydrogen

​

Bohr's ring model

Atomic emission Ground state vs excited state

• Ground State

• Lowest possible energy level of that e-

• Normal energy level of e-

• Excited State

  • Higher energy level

  • Caused by a gain of e-

  • Unstable and short-lived state

Atomic emission

• e- start in ground state

• e- absorb energy from source and jump to excited state

• e- release energy due to instability

  • Return to ground state

  • Released energy appears as light

Atomic emission spectrum

• Shows wavelengths of visible light released by excited electrons

• Unique for each element

  • Intensive property

• Viewed with a Spectroscope

  • Prism that separates colors of light

Photoelectric effect

• Heinrich Hertz- 1887

• Electrons are ejected from atoms when they absorb high-energy emr

• Electrons move at higher speed with shorter wavelength

  • Electrons not emitted if light energy below substance's threshold energy

  • Often required blue light/UV

  • Great intensity (brightness) releases more e-

  • Does not affect threshold or e- speed

• Used in solar cells, cameras, etc.

quantum

• Max Planck - 1900

• Minimum energy produced by wavelength of light

• Each wavelength produces different quantum of energy

• Light energy is always a multiple of that wavelength's quantum

• Energy is proportional to frequency of radiation

Photon

• Albert Einstein - 1905

• Particle of radiation

• Zero mass

• Carrying a quantum of energy for that wavelength

  • Light is a stream of photons

• Like drops of water from a hose

• Higher brightness = more photons

Photon energy

• E = hv

  • h - Planck's constant = 6.626 x 10^-34 J x s

    • Formula if you don't have wavelength: you know that c = lambda v, v = c/lambda, then you simply just plug it into E = hv --> E = h c/lambda

    • Higher E

  • Higher v and lower lambda

Energy example problems

What is the energy of a photon with a frequency of 2.65 x 10⁹ Hz?

E = hv    E = 6.626 x 10^-34(2.65 x 10^9) = 1.76 x 10^-24 J

What is the frequency of a photon with 3.45 x 10^-13 J of energy?

E = hv   3.45 x 10^-13 = 6.626 x 10^-34v --->  v = 3.45 x 10^-13/6.626 x 10^-34

5.21 x 10²⁰ Hz

Multiple Choice

You try! Find the energy of an infrared photon with a wavelength of 933 nm given the formula:

E = hv ---> E = hc/lambda

h = 6.626 x 10^-34 J x s

c = 3.00 x 10⁸

The hint is that you have to use 10^-9 for the 933 nm

• None of the above

• 2.13 x 10^-19 J

• 6.626 x 10^-34 J

• 2.13 x 10^-57 J

How we got that answer

E = hv ---> E = hc/lambda = 6.626 x 10^-34 x 3.00 x 10⁸/933 x 10^-9

2.13 x 10^-19 J

Ionizing Radiation

• EMR with enough energy to remove electrons from atoms

• Higher energy than visible light (high UV, X-ray, Gamma)

• Cause sunburns, cancerous cells, etc.

Non ionizing Radiation

• Low energy EMR

  • Single photons can't remove electrons

  • Some can excite electrons

  • Radio, Microwave, IR, Visible, low UV

• Can create heat when absorbed

Wave-particle duality

• Light is both a wave and a particle

• Louis de Broglie - 1924

• Matter has wave properties

  • Distinct wavelengths and frequencies

• Waves have particle properties

  • Photoelectric effect only possible if radiation is particles

De brogile wavelength formula

• Lambda = h/mv

  h = Planck's constant

  m = mass

  v = velocity

• AKA: Electron Cloud Model

• Erwin Schrodinger - 1926

• Described changes to system over time

• Schrodinger equation is solved to indicate probable regions where e- is located

​

Quantum mechanical model

Principle quantum number (n)

• Energy level of e-

  • aka Shell

• Determines size of area where e- can be found

• Higher n = larger area for movement

• Number from 1-7

Angular momentum quantum number (1)

• Shape of suborbital

• Shapes are bigger for higher energy levels

• Spherical

• 1 orbital per energy level

• 2 e- per energy level

​

S suborbitals

• Dumbbell-shaped

• 3 orbitals per energy level

• 6 e- per energy level

• n > 1

​

P suborbitals

• 5 orbitals per energy level

• 10 e- per energy level

• n > 2

​

D Suborbitals

• 7 orbitals per energy level

• 14 e- per energy level

• n > 3

​

F Suborbitals

Magnetic quantum number (ml)

• Which orbital the electron is in

• Gives the axis orientation (X,Y,Z)

Electron spin quantum number (ms)

• Which e- in orbital

• Electrons spin opposite each other

  • Cause magnetic attraction to lessen effect of electric repulsion

• Either +1/2 or -1/2

Where exactly are the electrons?

• Hard to tell

• Electrons are almost like spinning fans

• Uncertainty Principle

• Werner Heisenberg - 1927

• More you know about position of e-, less you know about where it's going

• Best to describe regions e- can be found, not exact locations

electron configuration

• Electrons fill in an atom in a specific order

• Each element has a unique configuration

• Configuration determines bonding properties and ratios

• There are no 5p-7p, 5d-7d or 5f-7f orbitals that are possible yet.

aufbau principle

• Electrons fill in orbitals with increasing energy

  • Ground State

• Suborbitals are completely filled before next lowest energy starts

• 4s has lower energy than 3d

• Out of order = excited state

  S block on periodic table rows 1-88 D block is rows 21-112 P block rows 5-118

  F block rows 58-103

Writing electron configuration examples

 

1. Hydrogen

H - 1 electron  so we start with the first block 1s 1, so you have 1 electron in a 1 s suborbital

2. Helium

He - 2 electrons, you can do s suborbital bc the max is 2  so 1s 2

Writing electron configuration examples con't

 

3. Lithium

Li - 3 electrons, so you can do 1s 2 and 2s 1

4. Carbon

C - 6 electrons, so you can do 1 s 2 (bc 2 out of 6 can go to s) 2s 2 (for the other 2) and then finally 2p 2 (Carbon is in the p block)

In general...here are the steps to writing electron configuration

1. Find the element you've been given on the periodic table.

2. See how many electrons that element has if you want (not rlly important)

3. Mark that element as your "ending point" as your starting point is going to be all the way at Hydrogen.

4. So in the Helium example, since Hydrogen is in the S block we can do 1s because it's in the 1st row and in the s block. Our goal is to get to Helium. Helium is also in the S block so we simply write 1s2. If we count, we successfuly make it to our ending point. The number after the row is the "steps" you take to get to that element.

Multiple Choice

Try this! Write the electron configuration for Argon.

S block on periodic table rows 1-88 D block is rows 21-112 P block rows 5-118

F block rows 58-103

• 5s p, 8s 6, 9d 5, 10s 6

• 1s 2, 2s 2, 3s 6, 4s 2, 5s 6

• Just leave me alone...

•  1s 2, 2s 2, 2p 6, 3s 2, 3p 6

How to get that answer

Our "ending goal" is Argon.
We start at Hydrogen. Hydrogen and Helium are in row 1, they are both in the s block, so we simply do 1s 2. (1 represents the row, 2 represents Hydrogen and Helium, the 2 steps we want to take). We're not at our "ending goal" yet so we go to the next row so 2s because we're still in the s block and once again we can take 2 steps as its in the same block so 2s 2. We are still in row 2 but are going to go to a different block, the p block, so we do 2p 6. Then we go to row 3, 3s 2 and finally we're still in row 3 just once again in the p block so 3p 6.

Orbital diagrams

• Show arrangement of electrons in orbitals

• Max of 2 e- per box

• Draw all boxes in suborbital (even if empty)

• The arrows going up or down depends on the

  suborbital. So as an example, the p

  suborbital can hold 3 orbitals per energy

  level so draw 3 arrows going up then 1

  going down.

Noble gas configurations

1. Place symbol of previous noble gas in brackets so ex: [Xe] 6s 2 5d 1 4f 6

2. Continue configuration after last e- of noble gas

 

• Ex: Nickel (atomic number 28)

• 1s 2, 2s 2, 2p 6, 3s 2, 3p 6, 4s 2, 3d 8 [Ar] (Argon)

• [Ar] 4s 2 3d 8

Noble gas configurations exAMPLES

For example, if you have Beryllium, the noble gas would be Helium (the one above it). Then you have Helium as your starting point and you go from there to Beryllium making the rest of the configuration simply [He] 2s 2 instead of 1s 2, 2s 2.

Multiple Choice

Try this! Write the noble gas configuration for Silicon.

S block on periodic table rows 1-88 D block is rows 21-112 P block rows 5-118

F block rows 58-103

• [Ar] 3s 2 3p 2

• [Ne] 3s 2 3p 2

• 1s² 2s² 2p⁶ 3s² 3p²

• None of the above

How we got that answer

The noble gas above Silicon is Neon, so we start at Neon [Ne] and go from there to Silicon. So 3s 2 because we are in the 3rd row, s suborbital and only 2 elements there correspond with that suborbital and then 3p 2, we are still in the 3rd row just going to the p suborbital and there's only 2 elements there to get to Silicon.

Lewis dot diagrams

• Indicate arrangement of valence e- in atom

• Draw valence e- as dots around chemical symbol

• Spread out in 4 directions --- MAX 8 dots

• Pair only when needed

Lewis dot diagram examples

For example, we have Carbon. Carbon has 4 valence electrons in total. So, we put the elements name so C in the middle and then one dot on the right, one below the C, one to the left and one above.

Draw

Draw the lewis dot diagram for Oxygen that has 6 valence electrons in total.

Your drawing should've looked like this:

The way you arranged the dots doesn't really matter as long as you have 2 dots that are unpaired and 2 that are.

If you didn't get that, here's how you should've approached the problem. First draw Oxygens symbol in the middle (O) then start by drawing 4 dots in total, 1 on the top, 1 on the bottom, 2 in total on the sides. Okay, so that's 4 out of 6. Then simply just pair the other two with whatever ones you want, so it can be on the sides maybe or on the bottom.

Atomic radius periodic trend

• Overall size of the atom

• Decreases from left to right across a period

  • Stronger nucleus, same e- energy level

• Increases from top to bottom in a group

  • More energy levels

• Larger change across period than down group

So the left side is bigger than the right in a row while for a column the bottom is bigger than the top.