General chemistry 1

Course notes for:

  • Course name: General Chemistry I
  • Course code: CHEM-1411
  • Instructor: Katheryn Townsend
  • Institution: South Plains College
  • Date: Spring 2022

Chapter 1 - The Study of Change

Chemistry is the study of matter and the changes it undergoes.

Classifications of matter

Matter is anything that occupies space and has mass.

  • pure substances
  • mixtures
  • elements
  • compounds
  • molecules
  • atoms

A pure substance is a form of matter that has a definite composition and distinct properties (i.e. water, oxygen, gold, sugar).

A mixture is a combination of two or more substances in which the substances retain their distinct identities. Physical means can be used to separate a mixture into its pure components.

  1. homogeneous - composition is the same throughout
  2. heterogeneous - composition is not the same throughout

An element is a substance that cannot be separated into simpler substances by physical or chemical means.

  • there are currently 118 identified elements
  • 98 occur naturally on Earth
  • 20 are synthetic
  • represented by one or two letter symbols. If two are used, the first is capitalized and the second is lowercase.

A compound is a pure substance composed of atoms of two or more different elements chemically bonded in fixed proportions. Compounds can only be separated into pure substances by chemical means.

A molecule is a pure substance composed of two or more atoms chemically bonded in fixed proportions. The atoms may be from the same element or different elements.

Atoms are the basic building blocks of matter.


A physical change does not alter the composition or identity of a substance.

H2O → H2O(l) → H2O(g)

A chemical change alters the composition or identity of the substance(s).

H2 + O2 = H2O

Classification of Matter by Physical State

Property Solid Liquid Gas
Shape fixed depends depends
Volume fixed fixed depends
Particle arrangement fixed random random
Particle closeness very close close far apart
Particle interaction very strong strong weak/none
Particle movement very slow/none moderate very fast
  • crystalline solids - orderly geometric patterns (i.e. salt, diamonds, graphite, sugar)
  • amorphous solids - no long-range pattern (i.e. plastic, gas, charcoal)

Change of state

Matter undergoes a change of state when it’s converted from one state to another. Every change in state requires a change in energy.

  • melting - solid to liquid
  • freezing - liquid to solid
  • sublimation - solid directly to gas
  • deposition - gas directly to solid
  • evaporation - liquid to gas
  • condensation - gas to liquid

A heating curve illustrates the changes of state as a solid is heated using sloped lines to show an increase in temperature. Plateaus indicate a change of state.

A cooling curve illustrates a change of state as a gas is cooled using sloped lines to show a decrease in temperature. Plateaus indicate a change of state.

Properties of matter

An extensive property of a material depends upon how much matter is being considered, i.e. mass, length, volume.

An intensive property of a material does not depend upon the amount of matter being considered, i.e. density, temperature, color.

International system of units (SI)

Name Symbol
Length meter m
Mass kilogram kg
Time second s
Electrical current ampere Amp
Temperature kelvin K
Amount mole mol
Luminous intensity candela cd

Metric prefixes

  • placed before base unit
  • gives the actual size of the measurement used
  • how many times the base unit is multiplied or divided by 10
King Henry Died By Drinking Chocolate Milk
kilo hecta decca base deci centi mili .. micro .. nano .. pico
k h da d c m μ n p

Derived units

Weight is the force that gravity exerts on an object. weight = c x mass

Volume: SI derived unit for volume is a cubic meter (m3). 1 L = 1000 ml = 1000 cm3 = 1 dm3 therefore, 1 ml = 1 cm3

Density compares the mass of an object to its volume. The SI derived unit for density is kg/m3. density = mass / volume

The density of pure water is 1g/mL

How does temperature affect density? Generally, heating an object causes if to expand. Therefore, the density will decrease.

Specific gravity (SPGR)

  • a comparison between densities
  • relationship between substance density and pure water density
  • unitless

SPGR = substance density / pure water density

Comparison of temperature scales

  • K = °C + 273.15
  • °C = K - 273.15
  • °F = °C * 1.8 + 32
  • °C = (°F - 32) / 1.8

Scientific notation


  • N is between 1-10
  • n is a positive or negative integer

Chapter 2 - Atoms, Molecules, and Ions

Periodic table


  • 7 horizontal rows
  • 7 periods
  • atomic number increases by 1
  • number of p+ increases by 1
  • number of e- increases by 1
  • properties change smoothly


  • vertical columns
  • group or family
  • number 1A-8A for main group = representative elements
    • 1A - alkali metals
    • 2A - alkali earth metals
    • 7A - halogens
    • 8A - noble gases
  • middle group = transition metals
    • lanthanides - elements number 58-71
    • actinides - elements number 90-103

3 divisions of the periodic table

  1. metals
  2. nonmetals
  3. metalloids
Metals Nonmetals Metalloids
largest group middle sized group smallest group
left side of staircase right side of staircase border the staircase (one exception)
good conductors not metals have properties that are intermediate between those of metals and nonmetals

Intro to bonding

A molecule is an aggregate of two or mote atoms in a definite arrangement held together by chemical forces.

  • a diatomic molecule contains only two atoms
  • a polyatomic molecule contains more than two atoms

Naturally occurring diatomic elements

  1. H
  2. N
  3. O
  4. F
  5. Cl
  6. Br
  7. I

An ion is an atom, or group of atoms, that has a net positive or negative charge.

  • cation - ion with a positive charge. If a neutral atom loses one or more electrons it becomes a cation.
  • anion - ion with a negative charge. If a neutral atom gains one or more electrons it becomes an anion.

A monatomic ion contains only one atom with a charge, i.e.:

  • Na+ - Sodium ion
  • Ca2+Calcium ion
  • Cl- - Chloride
  • O2- - Oxide

A polyatomic ion contains more than one atom with an overall charge, i.e.:

OH-, CN-, NH4+, NO3-

Why do atoms take on a charge when bonding?

So that the atom can become stable (not necessarily neutral) striving for the octet rule: an atom will either lose, gain, or share electrons in order to have 8 valence electrons. Valence electrons are the electrons on the outer most part of the atom.

  • group number = number of valence electrons
group 1A 2A 3A 4A 5A 6A 7A 8A
n valence electrons 1ve- 2ve- 3ve- 4ve- 5ve- 6ve- 7ve- 8ve-
gain +7e- +6e- +5e- +4e- +3e- +2e- +1e- +0e-
lose -1e- -2e- -3e- -4e- -5e- -6e- -7e- -0e-
charge +1 +2 +3 +3 -3 -2 -1 0
  • metals = +
  • nonmetals = -
  • C prefers to gain
  • Si prefers to gain
  • Ge prefers to lose
  • Sn acts like a transition metal
  • Pb acts like a transition metal
  • transition metals
    • always positive
    • more than one possibility

How do atoms bond?

Types of compounds:

  • ionic compounds
    • transfer of electrons
  • molecular (covalent) compounds
    • sharing of electrons
    • polar covalent - unequal sharing of e-
    • non-polar covalent - equal sharing of e-

Ionic compounds

Ionic compounds consist of a combination of cations and anions. The sum of the charges in each formula unit must equal zero.

  • cations = metals
  • anions = nonmetals

Chemical nomenclature

Ionic compounds:

  • metal + nonmetal
  • anion (nonmetal), add “ide” to element name
  • polyatomic names remain unchanged

Transition metal ionic compounds

  • indicate charge on metal with Roman numerals

Covalent compounds:

  • nonmetal + nonmetal
  • common names
  • element furthest left in periodic table is listed first
  • element closest to the bottom of group is listed first
  • EXCEPT: Hydrogen is placed between Nitrogen & Oxygen
  • use prefixes to indicate number of each kind of atom
  • last element ends in “ide” unless it’s polyatomic

Greek prefixes:

n prefix
1 mono
2 di
3 tri
4 tetra
5 penta
6 hexa
7 hepta
8 octa
9 nona
10 deca


An acid is a substance that yields hydrogen ions (H+) when dissolved in water.

  • HCL(s) - pure substance, hydrogen monochloride
  • HCLaq - Dissolved in water, hydrochloric acid

A simple acid is an acid that contains hydrogen (no oxygen) and another element. Has a hydro- prefix, an -ic suffix, and acid at the end. An oxoacid is an acid that contains hydrogen, oxygen, and another element.

Oxoanion Oxoacid
per-, -ate +1 O per-, -ic acid
-ate Representative Oxoanion (Polyatomic) -ic acid
-ite -1 O -ous acid
hypo-, -ite -2 O hypo-, -ous acid
-ide (simple anion) no O hydro-, -ic acid
  • only halogens have all five lines
  • everything else has only three lines: representative, 1 less oxygen, and no oxygens.

A base is a substance that yields hydroxide ions (OH-) when dissolved in water. Follows the same naming rule as simple anions.

Hydrates are compounds that have a specific number of water molecules attached to them. Anhydrous means “no water attached.” Name the compound in front of the water, then use Greek prefixes to indicate the number of water molecules.

Chapter 3 - Mass Relationships in Chemical Reactions

Atomic mass is the mass of an atom in atomic mass units (amu).

Molar mass is the mass of 1 mole of in grams.

The mole (mol) is the amount of a substance that contains as many elementary entities as there are atoms in exactly 12.00 grams of Carbon-12 (12C).

1mol=NA6.022x10231 \text{mol} = N_A \approx 6.022 \mathrm{x}{10}^{23}

where NA=N_A = Avogadro’s number

For any element: atomic mass (amu) = molar mass (g/mol).

Molecular mass (or molecular weight) is the sum of the atomic masses (in amu) in a molecule. For any element: molecular mass (amu) = molar mass (g/mol).

Formula mass is the sum of the atomic masses (in amu) in a formula unit of an ionic compound (metal + nonmetal). For any ionic compound: formula mass (amu) = molar mass (g/mol).

Conversion formulas:

gMMmol\frac{\text{g}}{\text{MM} \cdot \text{mol}}
  • g=MM×molg = \text{MM} \times \text{mol}
  • MM=gmol\text{MM} = \frac{g}{\text{mol}}
  • mol=gMM\text{mol} = \frac{g}{\text{MM}}
n atomsNAmol\frac{n \text{ atoms}}{N_A \cdot \text{mol}}
  • n atoms=NA×moln \text{ atoms} = N_A \times \text{mol}
  • NA=n atoms molN_A = \frac{n \text{ atoms }}{\text{mol}}
  • mol=n atoms NA\text{mol} = \frac{n \text{ atoms }}{N_A}

Percent composition of an element in a compound is (n x molar mass of element) / (molar mass of compound) x 100% where n is the number of moles of the element in 1 mole of the compound.

An empirical formula shows the simplest whole-number ratio of the atoms in a substance. A molecular formula shows the exact number of atoms of each element in the substance. Note: In the case of ionic compounds, the molecular formula and empirical formula are always the same.

Percent composition to empirical formula:

Mass percent (convert to grams; divide by molar mass) → moles of each element (divide by smallest number of moles) → mole ratios of elements (change to integer subscripts) → empirical formula

Balancing equations


  1. balance
  2. convert grams to moles
  3. use coefficients
  4. get into desired units

Reaction yield

  • Theoretical yield: the amount of product that should result at the end of the reaction (calculated)
  • Actual yield: the amount of a product obtained from a reaction (lab)

% yield = (actual / theoretical) x 100

Chapter 4 - Reactions in Aqueous Solutions

A solution is a homogeneous mixture of two or more substances.

The solute is(are) the substance(s) present in the smaller amount(s). It is the substance being dissolved.

The solvent is the substance present in the larger amount. It is the substance doing the dissolving.

When a solute dissolves, there are attractive forces between the solute particles holding them together (solute-solute interactions). There are also attractive forces between the solvent molecules (solvent-solvent interactions). When we mix the solute with the solvent, there are attractive forces between the solute particles and the solvent molecules (solute-solvent interaction). If the attractions between solute and solvent are strong enough, the solute will dissolve.

Salt vs. sugar dissolved in water:

  • salt: ionic compounds dissociate into ions when they dissolve.
  • sugar: molecular compounds do not dissociate when they dissolve.

An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.

Cations (+) and anions (-) are responsible for a solution’s electricity conductivity. A strong electrolyte has 100% dissociation. A weak electrolyte does not completely dissociate.

A nonelectrolyte is a substance that, when dissolved in water, results in a solution that does not conduct electricity.

No cations (+) or anions (-). The solution is all molecules.

A reversible reaction is a reaction that can occur in both directions.

compounds: {
  nonelectrolytes: [
    insoluble salts,
    covalent compounds,
    organic compounds,
  electrolytes: {
    strong: [
      strong acids (6),
      strong bases (10),
      soluble salts,
    weak: [
      weak acids,
      weak bases,

Types of reactions

  • precipitation reactions
  • acid and base reactions
  • oxidation - reduction reactions

Precipitation reactions

Reactions between aqueous solutions of ionic compounds that produce an ionic compound that is insoluble in water are called precipitation reaction, and the insoluble product is called a precipitate.

Writing ionic equations

  1. write the Name (word) equation.
  2. write the balanced molecular equation.
  3. write the ionic equation showing strong electrolytes completely dissociated into cations and anions.
  4. identify the spectator ions on both sides of the ionic equation.
  5. write the net ionic equation. Keep everything balanced and check charges.


  • must be aqueous
  • typically hydrogen plus a nonmetal
  • have a sour taste (vinegar, citrus fruits)
  • will cause color changes in plant dyes
  • will react with certain metals to produce hydrogen gas
  • will react with carbonates and bicarbonates to produce carbon dioxide gas
  • aqueous acid solutions conduct electricity

Strong acids:

  • HCl
  • HBr
  • HI
  • HNO2
  • H2SO4
  • HClO4


  • must be aqueous
  • typically a metal plus hydroxide
  • have a bitter taste
  • feel slippery (soap)
  • will cause color changes in plant dyes
  • aqueous base solutions conduct electricity

Strong bases:

  • 1A and 2A metals
  • except for Li and Be

Acid/base definitions

Arrhenius acid is a substance that produces H+1 (H3O+1) in water.

Arrhenius base is a substance that produces OH-1 in water.

A Bronsted acid is a proton donor.

A Bronsted base is a proton donor.

Neutralization reaction (acid-base reaction / H transfer)

acid + base → salt + water

Salt - a neutral ionic compound formed from an acid/base neutralization reaction.

Preparing solutions

Solubility is the maximum amount of a solute that dissolves in a specific amount of solvent.

A saturated solution contains the maximum amount of solute that can dissolve. There will usually be undissolved solute at the bottom of the container.

An unsaturated solution contains less than the maximum amount of solute. It can dissolve more solute.

The concentration of a solution is the amount of solute dissolved in a specific amount of solution.

soluteconcentrationsolution\frac{\text{solute}}{\text{concentration} \cdot \text{solution}}
  • solute = concentration x solution
  • concentration = solute / solution
  • solution = solute / concentration

Weight percent

Concentration by weight of solute in a solution.

weight percent = (g of solute / g of solution) x 100

Volume percent

volume percent = (mL of solute / mL of solution) x 100

Weight/volume percent

weight/volume percent = (g of solute / mL of solution) x 100


M = molarity = moles of solute / liters of solution


Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.

moles of solute before dilution (i) = moles of solute after dilution (f)

MiVi = MfVf

Solution stoichiometry and titrations

In a titration, a solution of accurately known concentration is gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.

The equivalence point is the point at which the reaction is complete.

The indicator is the substance that changes color at (or near) the equivalence point.

  1. balance the chemical equation
  2. get it into moles
  3. use the mole ratios
  4. get it into desired units

Oxidation-reduction reactions

  • oxidation half-reaction (loses electron)
  • reduction half-reaction (gains electron)

Oxidation number is the charge the atom would have in a molecule if electrons were completely transferred.

  • oxidation is loss of electrons
    • oxidized (reducing agent)
  • reduction is gain of electrons
    • reduced (oxidizing agent)

Electrons always go from the element being oxidized to the element being reduced.

Types of redox reactions

  • combination reaction: A + B = C
  • decomposition reaction: C = A + B
  • combustion reaction: A + oxygen = B
  • displacement reaction: A + BC = AC + B
  • disproportionation reaction: element is simultaneously oxidized and reduced

Activity series for halogens

F > Cl > Br > I

Chapter 5 - Gases

Physical characteristics of gases:

  • gases assume the volume and shape of their container
  • gases are the most compressible state of matter
  • gases will mix even and completely when confined to the same container
  • gases have much lower densities than liquid and solids

Kinetic molecular theory of gases

  1. A gas is composed of molecules that are separated from each other by distances far greater that their own dimensions. The molecules can be considered to be points; that is, they possess mass but have negligible volume.
  2. Gas molecules are in constant motion in random directions, and they frequently collide with one another. Collisions among molecules are perfectly elastic.
  3. Gas molecules exert neither attractive nor repulsive forces on one another.
  4. The average kinetic energy of the molecules is proportional to the temperature of the gas in kelvins. Any two gases at the same temperature will have the same average kinetic energy.

Kinetic energy and molecular velocities

  • Average kinetic energy of the gas molecules depends on the average mass and velocity: KE = 1/2 mv2
  • Gases in the dame container have the same temperature; therefore, they have the same average kinetic energy.
  • If they have different masses, the only way for them to have the same kinetic energy is to have a different…

Gas diffusion is the gradual mixing of molecules of one gas with molecules of another by virtue of their kinetic properties.

Molecular speed vs. molar mass

In order to have the same average kinetic energy, heavier molecules must have a slower average speed.

Temperature vs. molecular speed

As the absolute temperature (temperature in kelvin) increases, the average velocity increases.

Measurements to describe a gas

  • Pressure (P)
    • atm
    • mm Hg
    • torr
    • kPa
    • psi
  • Volume (V)
    • mL
    • L
  • Temperature (T)
    • K
  • Amount (n)
    • moles

Gases pushing

Pressure is the outward force created by collisions of gas molecules with the walls of the container.

The pressure exerted by a gas can cause some amazing and startling effects. Whenever there is a pressure difference, a gas will flow from areas of high pressure to low pressure.

Gas effusion is the process by which gas under pressure escapes from one compartment of a container to another by passing through a small opening.

Atmospheric pressure effects: differences in air pressure result in weather and wind patterns. The higher up in the atmosphere you climb, the lower the atmospheric pressure is around you and the lower the boiling point is for substances.

Pressure imbalance in ear: if there is a difference in pressure across the eardrum membrane, the membrane will be pushed out—what we commonly call a popped eardrum.

Instruments to measure pressure

  • Barometer: measures atmospheric pressure
  • Manometer: measures pressure of gases other than atmospheric

Common units of pressure

  • 1 atm = 760 mm Hg
  • 1 atm = 760 torr
  • 1 atm = 101.325 kPa
  • 1 atm = 14.7 psi

Gas laws

Relationships formed between P, V, T, and n.

  1. Boyle’s Law: the volume of a gas is inversely proportional to the pressure. P and V relationship. T and n remain constant. Decreasing the volume forces the molecules into a smaller space. More molecules will collide with the container, increasing the pressure.
  2. Charles’ Law: the volume of a gas is directly proportional to the absolute temperature (K). V and T relationship. P and n remain constant. At low temperatures, the gas molecules are not moving as fast; therefore, the volume is small. At high temperatures, the gas molecules are moving faster, causing the volume of the balloon to become larger.
  3. Gay-Lussac’s Law: T and P relationship. V and n remain constant.
  4. Avogadro’s Law: the volume of a gas is directly proportional to the number of gas molecules. V and n relationship. T and P remain constant. Increasing the number of gas molecules causes more of them to hit the wall at the same time. To keep the pressure constant, the volume must then increase.
  5. Combined Gas Law: combines Boyle’s, Charles’, and Gay-Lussac’s Laws. n remains constant.
  6. Ideal Gas Equation: combines all previous laws. R is the gas constant.
  7. Deviations from Ideal Behavior: two areas for corrections in the Ideal Gas Law (P and V).


  • The conditions 0 degrees Celsius and 1 atm are called standard temperature and pressure (STP)
  • Molar volume states that STP 1mol of a gas occupies a volume equal to 22.4 L.

Gas stoichiometry

Density calculations:

  • M is the mass of the gas in grams
  • MM is the molar mass of the gas

Molar mass of a gaseous substance:

  • d is the density of the gas in g/L

Mixtures of gases

When gases are mixed together, their molecules behave independently of each other. All gases in the mixture have the same volume and temperature. Therefore, in certain applications, the mixture can be thought of as one gas.

Partial pressure

The pressure of a single gas in a mixture of gases is called its partial pressure*. We can calculate the partial pressure if we know what fraction of the mixture it composes and the total pressure or the number of moles of the gas in a container of known volume and temperature. Dalton’s Law of Partial Pressures: The sum of the partial pressures of all the gases in the mixture equals the total pressure. P1 + P2 … = Ptotal

Mole fraction:

x1 = mole fraction of gas 1 = moles of gas 1 / total moles of gas

Chapter 5 formula list

  • Boyle’s: P1V1 = P2V2
  • Charles’: V1 / T1 = V2 / T2
  • Gay-Lussac’s: P1 / T1 = P2 / T2
  • Avogadro’s: V1 / n1 = V2 / n2
  • Combined: P1V1 / T1 = P2V2 / T2
  • Ideal: PV = nRT
  • STP: T = 0°C and P = 1atm
  • Molar volume: n = 1mol and V = 22.4L
  • R = 0.0821 L atm / mol K

Chapter 6 - Thermochemistry


Thermodynamics is the scientific study of the interconversion of energy.

Thermochemistry is the study of relationships between chemistry and energy. Even though chemistry is the study of matter, energy affects matter.

Energy is anything that has the capacity to do work.

Work is a force acting over a distance.

State functions are properties that are determined by the state of the system, regardless of how that condition was achieved. ΔE = Efinal - Einitial

Law of Conservation of Energy: Energy can be converted from one form to another or exchanged between objects, but cannot be created or destroyed.

ΔE = q + w ; where ΔE is the change in internal energy of a system, q is the head exchange between the system and the surroundings, and w is the work done on or by the system.

Exothermic process is any process that gives off heat—transfers thermal energy from the system to the surroundings.

Endothermic process is any process in which heat has to be supplied to the system from the surroundings.

The universe

The universe never loses or gains energy: ΔEuniverse = 0

universe = system + surroundings

The system is the specific part of the universe that is of interest in the study. The surroundings are everything else.

System types:

Type Exchange
open mass & energy
closed energy
isolated nothing

ΔEsystem + ΔEsurroundings = 0

ΔEsystem = -ΔEsurroundings

Chemical reaction

ΔEsystem = Efinal - Einitial

ΔEsystem = Eproduct - Ereactant

Chemical energy lost by combustion is energy gained by the system.

Signs of ΔEsys

When ΔEsys is (-) energy is given off by the system, exothermic When ΔEsys is (+) energy is absorbed by the system, endothermic

Sign conventions for work and heat

  • work done by the system on the surroundings: -
  • work done on the system by the surroundings: +
  • Head absorbed by the system from the surroundings (endothermic): +
  • Head absorbed by the surroundings from the system (exothermic): -


When a gas expands against a constant external pressure: w = -PΔV

Work is not a state function.


  • Pressure = atm
  • Volume = L
  • Work = J (Joule)

1 Liter-atmosphere = 101.3J


Enthalpy (H) is used to quantify the heat flow into of out of a system in a process that occurs at a constant pressure.

ΔH = Hproducts - Hreactants

if Hproducts < Hreactants then ΔH < 0 (exothermic)

if Hproducts > Hreactants then ΔH > 0 (endothermic)


Calorimetry is the measure of heat changes. A calorimeter is a closed container designed specifically for calorimetry.

The specific heat (s) of a substance is the amount of heat (q) required to raise the temperature of one gram of the substance by one degree Celsius.

The heat capacity (C) of a substance is the amount of heat (q) required to raise the temperature of a given quantity (m) of the substance by one degree Celsius.

C = m x s

Heat (q) absorbed or released:

  • q = C x Δt
  • q = m x s x Δt

Constant-volume calorimetry

Constant volume bomb calorimeter. Used to measure heat of combustion.

Constant-pressure calorimetry

Chapter 6 formula list


  • ΔEsystem = -ΔEsurroundings
  • ΔE = Efinal - Einitial
  • ΔE = Eproduct - Ereactant
  • constant pressure
    • w = -PΔV (101.3J / 1 Latm)
    • q = ΔH
  • ΔE = q + w
    • or: ΔE = ΔH - PΔV
  • ΔE = ΔH - RT(Δn)
  • R = 8.314J/mol K
  • Δn = #moles of product gasses - #moles of reactant gases


  • C = m x s
  • q = C x Δt or q = m x s x Δt
  • constant volume: qrxn = -CΔt
  • constant pressure: qrxn = - (qwater + qcal)

Chapter 7 - The quantum-mechanical model of the atom

Electron behavior determines much of the behavior of atoms. Directly observing electrons in the atom is impossible—the electron is so small that observing it changes its behavior.

The quantum-mechanical model explains how electrons exist and behave in atoms. It helps us understand and predict the properties of atoms that are directly related to the behavior of the electrons.

The nature of light

Light is a form of electromagnetic radiation composed of perpendicular oscillating waves: one for the electric field and one for the magnetic field. An electric field is a region where an electrically charged particle experiences a force. A magnetic field is a region where a magnetized particle experiences a force.

A wave is a vibrating disturbance by which energy is transmitted. Wavelength (λ) is the distance between identical points on successive waves. Amplitude is the vertical distance from the midline of a wave to the peak or trough. Frequency (ν) is the number of waves that pass through a particular point in one second (Hz = 1 cycle/s).

The speed of the wave = wavelength x frequency:

  • u = λν

Electromagnetic radiation is the emission and transmission of energy in the form of electromagnetic waves. All electromagnetic waves have the same constant speed: speed of light (c) in vacuum = 3 x 108 m/s

For all electromagnetic radiation, wavelength x frequency = the speed of light:

  • λν = c


White light is a mixture of all the colors of visible light. When an object absorbs some of the wavelengths of white light while reflecting others, it appears colored. The color of light is determined by its wavelength. The brightness is determined by its amplitude.

The electromagnetic spectrum

Visible light comprises only a small fraction of all the wavelengths of light called the electromagnetic spectrum. Short wavelength (high frequency) light has high energy (Gamma ray is the highest energy). Long wavelength (low frequency) light has low energy (Radio wave is the lowest energy).

High to low energy waves:

  1. gamma rays
  2. x rays
  3. ultraviolet
  4. visible
  5. infrared
  6. microwave
  7. radio wave


The interaction between waves is called interference. When waves interact so that they add to make a larger wave, it is called constructive interference (in phase). When waves interact so they reduce or cancel each other, it is called destructive interference (out of phase).


When traveling waves encounter an obstacle or opening in a barrier that is about the same size as the wavelength, they bend around it. This is called diffraction. Traveling particles do not diffract.

Planck’s quantum theory

Planck noticed that when solids are heated they give off electromagnetic radiation over a wide range of wavelengths. He observed that the energy was not being emitted (or absorbed) continuously but instead in little bundles he called quanta. He concluded that the energy given off by light is proportional to the frequency. The constant that relates them is Plank’s constant (h): 6.63 x 10-34 J*s. Planck called this amount of energy quantum. E = hν

Einstein’s photoelectric effect

Photoelectric effect - electrons are ejected from the surface of certain metals exposed to light at least a certain minimum frequency, the threshold frequency. The behavior of these ejected electrons did not follow wave theory so Einstein proposed that the beam of light is actually a stream of particles. Einstein called these particles photoelectrons and every photon has energy.

Einstein’s work paved the way for the solution to the emission (light) spectra of atoms. emission spectra - either continuous or line spectra of radiation emitted by substances.

Bohr’s theory of the hydrogen ion

Bohr’s model of the atom included the idea of electrons moving in circular orbits, but he imposed that the single electron in the hydrogen atom could be located only in certain orbits. Because each orbit has a particular energy associated with is, the energies associated with the electron motion must be fixed in value, or quantized. As an electron drops from a higher energy orbit to a lower one there is an emission of radiation, giving off energy in the form of light. En = -RH (1 / n2) where R (Rydberg constant) = 2.18 x 10-18J and n is the shell number.

  • Ephoton = ΔE = Ef - Ei
    • Ef = -RH (1 / n2f)
    • Ei = -RH (1 / n2i)


  • ΔE = -RH(1/n2f - 1/n2i)

Schrodinger wave equation

Three values represent the size, shape, and orientation of an orbital in an atom. A fourth value represents the spin of the electron located in that orbital.

Quantum numbers:

  1. principal quantum number: n
  2. angular momentum quantum number: l
  3. magnetic quantum number: ml
  4. magnetic spin quantum number: ms

Ψ = fn(n, l, ml, ms)

  • shell - electrons with the same value of n
  • subshell - electrons with the same values of n and l
  • orbital - electrons with the same value of n l and ml

Pauli exclusion principle - no two electrons in an atom can have the same four quantum numbers.

Principal quantum number

n is how close the electron is to the nucleus.

Angular momentum

l is a range of whole numbers from 0 to n-1 telling the shape of the orbital.

  • l = 0: s orbital = 1 orbital
  • l = 1: p subshell = 3 orbitals
  • l = 2: d subshell = 5 orbitals
  • l = 3: f subshell = 7 orbitals

Magnetic quantum number

ml is a range of whole numbers from -l to +l telling the orientation of the orbital in space.

Magnetic spin quantum number

Each orbital can hold a maximum of two electrons represented by ms with a value of either +1/2 or -1/2. Positive is spinning with the magnetic field and negative is spinning against.

Quantum number tables

n = 1 n = 2 n = 3 n = 4
l = 0(s) l = 0(s), 1(p) l = 0(s), 1(p), 2(d) l = 0(s), 1(p), 2(d), 3(f)
l = 0(s) l = 0(s), 1(p) l = 0(s), 1(p), 2(d) l = 0(s), 1(p), 2(d), 3(f)
ml = 0 ml = -1, 0, 1 ml = -2, -1, 0, 1, 2 ml = -3, -2, -1, 0, 1, 2, 3
# orbitals in a shell max # electrons in a shell
n2 2n2

Electron configuration

Electron configuration is how the electrons are distributed among the various atomic orbitals in an atom. Formatted as 1s1 (nlnum electrons).

“Fill up” electrons in lowest energy orbitals first (Aufbau principle).

Orbital diagram

The most stable arrangement of electrons in subshells is the one with the greatest number of parallel spins (Hund’s rule).

  • unpaired electrons makes an element paramagnetic
  • all paired electrons makes an element diamagnetic

Chapter 8 - Periodic relationships among the elements

Periodic law: when the elements are arranged in order of increasing atomic mass, certain sets of properties recur periodically. Elements with similar properties were placed in the same column.

Mendeleev’s periodic law allows us to predict what the properties of an element will be based on its position on the table. Quantum mechanics is a theory that explains why the periodic trends in the properties exists.

Trends gradually change across a period of down a group. There are a few diagonal relationships. There are a few exceptions.

isoelectronic: ions and/or atoms that have the same number of electrons and the same electron configuration.

Representative elements cations/anions

Transition metals cations/anions

When a cation is formed from an atom of a transition metal, electrons are always removed first from the ns orbital and then from the (n - 1)d orbitals.


In a multielectron system, electrons are simultaneously attracted to the nucleus and repelled by each other. Outer electrons are shielded from the nucleus by the core electrons. Because of this shielding, the outer electrons do not experience the full strength of the nuclear charge.

Effective nuclear charge (Zeff) is the positive charge felt by an electron.

Zeff = Z - S

  • Zeff is the effective nuclear charge, or Z effective
  • Z is the number of protons in the nucleus, the atomic number
  • S is the average amount of electron density between the nucleus and the electron

Atomic radius

atomic radius is half the distance between two nuclei in two adjacent atoms. R = d/2

On the periodic table, atomic radius increases from right to left and top to bottom. Except, transition metals, however, increase in size down the group but stay roughly the same size across the d block.

  • Cation is always smaller than its neutral atom
  • Anion is always larger that its neutral atom

Ionization energy

Ionization energy is the minimum energy (kJ/mol) required to remove an electron from a gaseous atom in its ground state.

Trend: I1 < I2 < I3

Ionization energy increases from let to right and bottom to top (periodic table). Except, from 2A to 3A, 5A to 6A.

Electron affinity

Electron affinity is the negative of the energy released when a neutral atom gains an electron.

EA decreases top to bottom (periodic table). Except, increase from the second period to the third period. EA increases across a period.

Properties of metals and nonmetals


  • malleable and ductile
  • shiny, lustrous, reflect light
  • conduct heat and electricity
  • most oxides are basic
  • form cations
  • form ionic compounds
  • lose electrons in reactions (oxidized)


  • brittle in solid state
  • dull, non-reflective solid surface
  • electrical and thermal insulators
  • most oxides are acidic
  • form anions and polyatomic anions
  • form covalent compounds
  • gain electrons in reactions (reduced)

Metallic character

Metallic character is how closely an element’s properties match the ideal properties of a metal. More malleable and ductile, better conductors, and easier to ionize. Metallic character decreases from left to right and increases from top to bottom (periodic table).

Chapter 9 - Chemical bonding I

Types of chemical bonds:

  • ionic
  • covalent
    • polar
    • non-polar
  • metallic

Electronegativity is the tendency for an atom of a given chemical element to attract shared electrons when forming a chemical bond.

Chapter 10 - Chemical bonding II

Lewis electron dot structure are diagrams that show the bonding between atoms of a molecule, as well as the lone pairs of electrons that may exist in the molecule.

Formal charge

The formal charge of an atom is computed as the difference between the number of valence electrons that a neutral atom would have and the number of electrons that belong to it in the Lewis structure.


For some molecules and ions, it is difficult to determine which lone pairs should be moved to form double or triple bonds, and two or more different resonance structures may be written for the same molecule or ion. In such cases it is usual to write all of them with two-way arrows in between. When this situation occurs, the molecule’s Lewis structure is said to be a resonance structure, and the molecule exists as a resonance hybrid.