General Chemistry final Vocab

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Chemistry Final Exam Study Guide:

Chapter 1:

Chemistry- the study of matter and its interactions

Theory- a confirmed and accepted explanation of the laws of nature

Hypothesis- a possible explanation for an observation

Law- when a statement can summarize a large number of observations

Atoms- the building blocks of matter

Element- a substance that cannot be decomposed into a simpler substance by normal chemical means

Substance- a material that is chemically the same throughout

Compound- a substance that can be decomposed into its elements by chemical processes

Mixture- a combination of two or more substances that can be separated by differences in the physical properties of the substance

Property- anything we observe or measure about a sample of matter

Physical Properties- one that can be observed without changing the substances present in the sample

Ex: Mass, volume, color, phase, boiling point

Chemical Properties- the tendency to react to form new substances

Ex: explosiveness, flammability, corrosiveness

Physical Change- occurs without a change in the composition of the substance

Ex: Melting of ice, freezing

Chemical Change- occurs when the substance is converted to a different type of matter

Ex: rusting of iron, burning of wood

Extensive properties- those that depend on the size of the sample; they measure how much matter is in a particular sample

Ex: mass, volume

Intensive Properties- those that are independent of the size of the sample; they depend on what the sample is

Ex: color, melting point, density

Heterogeneous Mixture- a mixture in which different parts have different properties and comp

Ex: blood, chicken noodle soup

Homogeneous Mixture (solution)- a mixture in which all parts of the sample exhibit identical properties

Ex: sugar water, milk, air

Cation- when an atom loses an electron (Ca2+)

Anion- when an atom gains an electron (P3-)

Chapter 2:

Law of Conservation of Mass- there is no detectable change in mass when a chemical reaction occurs

Law of Constant Composition- all samples of a pure substance contain the same elements in the same properties by mass

Law of Multiple Proportions: in each pair of compounds formed by the same elements, the masses of one element that combine with a fixed mass of a second element are always in a ratio of small whole numbers

Dalton’s Atomic Theory-

Matter is composed of small indivisible particles called atoms.

An element is composed of entirely one type of atom. The chemical properties of all atoms of any element are the same.

A compound contains atoms of two or more different elements. The relative number of atoms of each element in a particular compound is always the same.

Atoms do not change their identities in chemical reactions. Chemical reactions rearrange only how atoms are joined together.

The Electron Discovery-

Streams of negatively charged particles were found to emanate from cathode tubes, causing fluorescence

J.J. Thomson credited with discovery in 1897

Millikan Oil-Drop Experiment

Robert Millikan determined the charge on the electron (negative) in 1909

Millikan injected tiny droplets of oil into a chamber and exposed the chamber contents to high-energy radiation, causing each oil drop to acquire an electrical charge.

Radioactivity- the spontaneous emission of radiation by an atom

The Rutherford Experiment-

Discovery of the nucleus

“Nuclear Model of the Atom”

Ernest Rutherford shot alpha particles at a thin sheet of gold foil and observed the pattern of scatter of the particles

Rutherford postulated a very small, dense nucleus with the electrons around the outside of the atom

Protons were discovered by Rutherford in 1919

Neutrons were discovered by James Chadwick in 1932

Proton- positive charge

Electron- negative charge

Neutron- no charge

The nucleus contains protons and neutrons, electrons surround the nucleus

Mass number- the number of protons plus neutrons

Isotopes- atoms of the same element with different masses and different numbers of neutrons

Molecular compounds- composed of molecules and almost always contain nonmetals

Ionic Compounds- generally formed between metals and nonmetals

Organic Chemistry- the study of carbon

Hydrocarbons- organic compounds that contain only the elements hydrogen and carbon

Dissociation- the separation of a compound into smaller units

Ion- formed by either adding or subtracting electrons from the atom

Chapter 3:

Chemical equations- concise representations of chemical reactions

Reactants- appear on the left side of the equation

Products- appear on the right side of the equation

Reaction Types:

Combination reaction- two or more substances react to form one product

2Mg(s) + O2(g) 2MgO(s)

N2(g) + 3H2(g) 2NH3(g)

C3H6(g) + Br2(l) C3H6Br2(l)

Decomposition Reaction- one substance breaks down into two or more substances

CaCO3(s) CaO(s) + CO2(g)

2KClO3(s) 2KCl(s) + O2(g)

2NaN3(s) 2Na(s) + 3N2(g)

Combustion reaction- generally rapid reactions that produce a flame, result in carbon dioxide and water

CH4(g) + 2O2(g) CO2(g) + 2H2O(g)

C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(g)

Oxidation-Reduction Reaction- electrons are transferred from one species to another. Typically include something in its elemental form.

Cu(s) + 2AgNO3(aq) 2Ag(s) + Cu(NO3)2(aq)

Avogadro’s Number- 6.022 x 1023

Use Avogadro’s number when finding the number of atoms of a substance

Limiting Reactant- the reactant present in the smallest stoichiometric amount

Chapter 4:

Solutions- homogeneous mixtures of two or more pure substances

Solvent- present in greatest abundance

Solute- all other substances

Dissociation- when an ionic substance dissolves in water, the solvent pulls the individual ions from the crystal and solvates them

Electrolyte- a substance that dissociates ions when dissolved in water

Nonelectrolyte- may dissolve in water, but it does not dissociate into ions when it does so

Strong electrolyte- dissociates completely when dissolved in water (strong acids, strong bases, soluble ionic salts)

Weak electrolyte- only dissociates partially when dissolved in water

Strong Acids:

Hydrochloric (HCl)

Hydrobromic ((HBr)

Hydroiodic (HI)

Nitric (HNO3)

Sulfuric (H2SO4)

Chloric (HClO3)

Perchloric (HClO4)

Strong Bases- soluble metal salts of hydroxide ion containing: alkali metals, calcium, strontium, barium

Group 1A Metal Hydroxides

LiOH

NaOH

KOH

RbOH

CsOh

Heavy Group 2A Metal Hydroxides

Ca(OH)2

Sr(OH)2

Ba(OH)2

Precipitation Reaction- when one mixes ions that form compounds that are insoluble

Molecular equation- lists reactants and products in their molecular form

Ionic equation- all strong electrolytes are dissociated into their ions

Steps to writing Net Ionic Equations:

Write a balanced molecular equation

Dissociate all strong electrolytes

Cross out “spectator ions”

Write net ionic equation with remaining species

Acid-Base Reaction- acid donates (H+) to the base

Neutralization reaction- solutions of an acid and a base are combined the products are a salt and water

Gas-Forming Reaction- the expected product decomposes to give a gaseous product

Oxidation- occurs when an atom or ion loses electrons

Reduction- occurs when an atom or ion gains electrons

Displacement reactions- ions oxidize element then ions are reduced

Chapter 5:

Energy- the ability to do work or transfer heat

Kinetic Energy- energy an object possesses by virtue of its motion

Potential energy- energy an object possesses by virtue of its position or chemical composition

The SI unit of energy is the Joule (J)

Work- energy used to move an object some distance

Exergonic- when energy is released

Endergonic- when energy is absorbed

Endothermic- when heat is absorbed by the system from the surroundings

Exothermic- when heat is released by the system into the surroundings

Enthalpy- the internal energy plus the product of pressure and volume

Heat capacity- the amount of energy required to raise the temperature of a substance by 1 K or 1 C

Specific Heat Capacity- the amount of energy required to raise the temperature of 1 g of a substance by 1 K or 1 C

Hess’s Law- “if a reaction is carried out in a series of steps, the change in enthalpy for the overall reaction will be equal to the sum of the enthalpy changed for the individual steps”

Chapter 6:

Pressure- the amount of force applied to an area

Atmospheric Pressure- the weight of air per unit of area

Units of Pressure Conversions:

1 atm =

101.325 kPa

1.01325 bar

760 torr

760 mmHg

14.69 psi

Standard Pressure- normal atmospheric pressure at sea level (1 atm)

Boyle’s Law- the volume of a fixed quantity of gas at constant pressure is inversely proportional to the pressure

Charles’s Law- the volume of a fixed amount of gas at constant pressure is directly proportional to its absolute temperature

Avogadro’s Law- the volume of a gas at constant temperature and pressure is directly proportional to the number of moles of the gas

Ideal Gas Equation Constant- R

Dalton’s Law of Partial Pressures- the total pressure of a mixture of gases equals the sum of the pressures that each would exert if it were present alone

Effusion- the escape of gas molecules through a tiny hole into an evacuated space

Diffusion- the spread of one substance throughout a space or throughout a second substance

Chapter 7:

Wavelength- the distance between corresponding points on adjacent waves

Frequency- the number of waves passing a given point per unit of time

Principal Quantum Number (n)- describes the energy level on which the orbital resides

Angular Momentum Quantum Number- defines the shape of the orbital

Magnetic Quantum Number- describes the three-dimensional orientation of the orbital

Shell- orbitals with the same value of n

Subshell- different orbital types within a shell

Degenerate- orbitals on the same energy level have the same energy

Hund’s Rule- “for degenerate orbitals, the lowest energy is attained when the number of electrons with the same spin is maximized.”

Chapter 8:

Ionic Bond- electrostatic attraction between ions

Covalent Bond- sharing of electrons

Metallic- metal atoms bonded to several other atoms

Lattice energy- the energy required to completely separate a mole of a solid ionic compound into its gaseous ions

Electronegativity- the ability of atoms in a molecule to attract electrons to themselves

Bond dipole- when two atoms share electrons unequally

Writing Lewis Structures:

Find the sum of valence electrons of all atoms in the polyatomic ion or molecule

If it is an anion, add one electron for each negative charge

If it is a cation, subtract one electron for each positive charge

The central atom is the least electronegative element that isn’t hydrogen. Connect the outer atoms to it by single bonds

Fill the octets of the outer atoms

Fill the octet of the central atom

If you run out of electrons before the central atom has an octet, form multiple bonds until it does

Formula Sheet:

Temperature Conversions:

K = °C + 273.15

°C = 5/9 (°F - 32°)

°F = 9/5 (°C + 32°)

Density:

D = M/V

Molarity:

M= mol/L

Ideal Gas Equation:

PV = nRT

P = pressure (atm)

V = volume (L)

n = moles

R (constant) = 0.08206 (L*atm/mol*K)

T= temperature (K)

Avogadro’s Number (constant):

6.022*1023

Solutions & Dilutions:

m = C * vol * MW

m = mass (g)

C = molarity (M)

CiVi = CfVf

Finding molar mass of a gas:

M = dRT/P

Finding the density of a gas:

D = PMW/RT

Finding Change in P, V, or T of a gas:

P1V1/T1 = P2V2/T2

Planck’s Constant (h):

6.626*10-34 J-s

Speed of light constant (c):

3.0*108 m/s

Converting nm to m:

1 nm = 1*10-9 m

1 m = 1*109 nm

Light as a wave:

λ = C/ν

λ = wavelength (m)

ν = frequency (Hz or s-1)

c = speed of light constant

Light as a photon:

ν = E/h

E = energy (J)

h = planck’s constant

Matter as a wave:

λ = h/mv

λ = wavelength (m)

h = planck’s constant

m = mass (kg)

v = velocity