NC State Chemistry Placement Topics

Stoichiometry

deals with the relationships between the elements making up substances and the property of the substances

Law of Definite Proportions

deals with chemical compounds, or substances that consist of two or more elements; states that the element of every pure compound exist in the same ratio by mass

Law of Multiple Proportions

states that when chemical elements combine in a chemical reaction to form a compound, they combine in a ratio of small whole numbers

Avagadro's number

one mole (mol) = 6.02 x 10^23

Periodic Table

Periodic Table Trends

Moving Left to Right: atomic radius decreases, ionization energy increases, electronegativity increases.

Moving TOP to BOTTOM: atomic radius increases, ionization energy decreases, electronegativity increases

Gas Laws

experiments with a large number of gases reveal that the state, or condition, of many gaseous substances can be defined using four variables: temperature (T), pressure (P), volume (V), and the quality

Mass

Kilogram (kg)

Length

Meter (m)

Time

Second (s or sec)

Temperature

Kelvin (K)

Amount of substance

Mole (mol)

Electric current

Ampere (A or amp)

luminous intensity

Candela (cd)

Peta (P)

10^15

Tera (T)

10^12

Giga (G)

10^9

Mega (M)

10^6

Kilo (k)

10^3

Deci (d)

10^-1

Centi (c)

10^-2

Milli (m)

10^-3

Micro (μ^b)

10^-6

Nano (n)

10^-9

Pico (p)

10^-12

Femto (f)

10^-15

Atto (a)

10^-18

Zepto (z)

10^-21

SI Units

Boyle's Law

states that for a fixed quantity of gas maintained at constant temperature, the volume of the gas is inversely proportional to the pressure; the product of the pressure (P) multiplied by the volume (V) remains constant if there is no change in temperature or in the number of particles inside the container; the pressure doubles when a gas is compressed to half its volume at constant temperature

PV = constant

Charles's Law

states that for a fixed quantity of gas at constant pressure, the volume of gas is directly proportional to its temperature; the ratio between the volume (V) of a gas and its temperature (T) remains constant if the pressure does not change; doubling the temperature doubles the gas's volume if the pressure does not change

V/T = constant

Avogadro's Law

states that at constant temperature and pressure, the volume of the gas is directly proportional to the moles of gas; equal values of different gases all contain the same number of particles if they all have the same pressure and temperature

Ideal Gas Law

PV = nRT

(P - pressure)

(V - volume)

(n - # of moles of gas)

(T - absolute temperature)

(R - universal gas constant - 8.314 joules per kelvin per mole)

Ways gas pressure can be doubled:

1. Gas can be squeezed into 1/2 its original volume

2. Twice as much gas can be forced into the original volume

3. Absolute temperature can be doubled

Solution

A homogeneous mixture of two or more individual substances; occurs when a substance dissolves in another; consists of solutes and solvents

Solute

substance that is dissolved

Solvent

substance that causes another substance to dissolve

Concentration of a solution

the amount of solute present in a given amount of solution

Concentrated solution

relatively large quantity of solute per unit amount of solution

Dilute solution

relatively small quantity of solute per unit amount of solution

Solubility

measure of how much solute dissolves in a given amount if solvent at a given temperature

Percent by weight

the percentage of mass of a component of a solute in a given mass of the solution

% by weight of solute = (grams solute/grams solution) ⋅ 100

Mole fraction (X)

ratio of the number of moles of a component to the total number of moles of all the components

X = (moles component/moles all components)

Molarity (M)

number of moles of solute in a liter of solution

M = (moles solute/liters solution)

Molality (m) of a solution

number of moles of solute in a kilogram of solvent

Mixture

two or more substances mixed together but not chemically combined

Homogeneous mixture

uniform throughout, like a solution

(ex: sugar and water)

Heterogeneous mixture

consists of two or more easily identifiable substances

(ex: sand and water)

As temperature increases

it's solubility increases for most solids and decreases for most gases

As pressure increases

the gases solubility increases

Saturation

the point at which a solution reaches a point where no more solute can be dissolved in it at the same temperature and pressure

If the solute-solvent attraction is stronger than the solvent-solvent attraction

the solute will dissolve readily

If the solute-solvent forces are weaker than the other forces

only a relatively small amount of the solute will dissolve

Solution forming process

1. Separation of solute molecules

2. Separation of solvent particles

3. Mixing of solvent and solute molecules so that the solute particles occupy positions that are normally taken by solvent molecules

Crystallization

process by which matter forms crystals, can occur once a solution has passed the saturation point

Units of temperature and conversion

Phase changes

at certain combinations of temperature and pressure, ice changes directly to steam without becoming a liquid first

Endothermic and exothermic reactions

Endothermic: changes in which energy is absorbed

Exothermic: changes in which energy is released

1- Ions

Acetate (CH3COO)

Benzoate (C6H5COO)

Chlorate (ClO3)

Chlorite (ClO2)

Cyanide (CN)

Dihydrogen Phosphate (H2PO4)

Glutamate (C5H8NO4)

Hydrogen Carbonate/Bicarbonate (HCO3)

Hydrogen Oxalate (HOOCCOO)

Hydrogen Sulfate/Bisulfate (HSO4)

Hydrogen Sulfide/Bisulfide (HS)

Hydrogen Sulfite/Bisulfite (HSO3)

Hydroxide (OH)

Hypochlorite (CLO) or (OCl)

Nitrate (NO3)

Nitrite (NO2)

Perchlorate (CLO4)

Permanganate (MnO4)

Stearate (C12H35COO)

Thiocyanate (SCN)

2- Ions

Carbonate (CO3)

Chromate (CrO4)

Dichromate (Cr2O7)

Hydrogen Phosphate (HPO4)

Oxalate (OOCCOO)

Silicate (SiO3)

Sulfate (SO4)

Sulfite (SO3)

Tetraborate (B4O7)

Thiosulfate (S2O3)

3- Ions

Borate (BO3)

Citrate (C3H4OH(COO)3)

Phosphate (PO4)

5- Ion

Tripolyphosphate (P3O10)

1+ Ions

Ammonium (NH4)

Hydronium (H3O)

2+ Ion

Mercury/l (Hg2)

The Dalton Model

described atoms as small, spherical, indivisible particles; theory states that each element is composed of its own kind of atoms, all with the same relative weight; explained why a fixed weight of one substance always combines with a fixed weight of another substance in forming a compound

Thompson Model

described the atom as consisting of smaller subatomic particles; discovered the first subatomic particle: the electron; determined that since all atoms are electrically neutral, the atoms must contain as many positive charges as they do negatively charged electrons

Rutherford Model

stated that the atom is mostly empty space; proposed the theory that the atom consists of a nucleus with a positive charge with enough electrons rotating around it to balance the charge; these electrons were kept within the atom by the attraction of the positively charged nucleus

Bohr Model

proposed that an atom's electrons could travel only in certain successively larger orbits around the nucleus; through the outer orbits could hold more electrons than the inner ones; suggested that the electrons in the outermost orbit determined the atom's chemical properties

Quantum Mechanics

explains the structure of atoms, the way they give off light, and other related matters

Principal Quantum Number

n

Higher quantum number

higher orbital energies

Ground state

lowest energy level; where hydrogen electrons were normally

Excited state

when an electron becomes unstable due to the absorption of energy

Acids and bases

two types of chemical compounds characterized by their opposite effects in certain physical and chemical properties and their ability to neutralize each other in a chemical reaction

Acids

occur naturally and some are essential for life; many acids are poisonous and strong acids can cause severe burns

(ex: Hydrochloric Acid (HCl) produced in the stomach to aid in digestion)

Bases

have many practical uses

(ex: Magnesium Hydroxide (Mg(OH)2) is often used as an ingredient in antacids)

Arrhenius Theory

observed that all substances classified as acids contain hydrogen ions, H+, and that all substances known as bases contain the hydroxide ions, OH-

Arrhenius acid

substance whose water solution contains a high concentration of hydrogen ions

Arrhenius base

substance whose water solution contains a high concentration of hydroxide ions

Acid properties

- have a sour taste and produce a prickling or burning sensation if they come into contact with the skin

- dissolve many metals

- turn blue litmus paper red

- are neutralized by bases

Base properties

- react with an acid to decrease or neutralize its acidic properties

- also called alkalis

- turn red litmus paper blue

- feel slippery and taste bitter when dissolved in water

Bronsted-Lowry Theory

theorized that an acid-base reaction is a proton-transfer reaction in which a proton (a hydrogen ion) is transferred from the acid to the base

theorized that the strength or weakness of different acids and bases is a measure of their tendency to lose or gain protons

Bronsted acid

proton (H+) donor

Bronsted base

proton acceptor

Bronstead-Lowry reaction

B + HA < ⇌ > HB+ A-

(B = base proton receiver)

(HA = acid proton receiver)

(A- = base)

(HB+ = acid)

Conjugate acid-base pairs

when combinations such as the acid HA and the base A- result from an acid losing a proton or a base gaining one

Acid: Hydrochloric Acid (HCl)

Base: Chloride ion (Cl-)

Acid: Nitric Acid (HNO3)

Base: Nitrate ion (NO3-)

Acid: Hydrocyanic Acid (HCN)

Base: Cyanide ion (CN-)

Acid: Perchloric Acid (HClO4)

Base: Perchlorate ion (ClO4-)

Acid: Sulfuric Acid (H2SO4)

Base: Hydrogen Sulfate ion (HSO4-)

A strong acid

donates protons easily

A weak acid

clings to its protons

A strong base

has a strong attraction for protons

A weak base

has a weak attraction for protons

The stronger the acid

the weaker its conjugate base

The weaker the acid

the stronger its conjugate base

Chemical Formulae

Compounds

substances that contain more than one kind of atom

Also includes:

- Hydrogen Peroxide (H2O2)

- Glucose (C6H12O6)

- Sucrose (C12H22O11)

- Propane (C3H8)

- Octane (C8H18)

- Methanol (CH3OH)

- Ethanol (C2H5OH)

Molecular formulas

show the exact number and type of atoms combined in each molecule of a compound

(ex: Methane, CH4, contains one carbon atom and four hydrogen atoms)

Empirical formulas

show the simplest ratios of the ions--atoms or molecules that have an electric charge--present in an ionic compound

(ex: Sodium Chloride is NaCl, while Barium Fluoride is BaF2)

Ionic compounds do not form individual, discrete molecules but instead are arranged in a crystal lattice; their empirical formula must reflect an overall net charge of zero