Chemistry - All Vocabulary

Chemistry

the study of matter and how matter works with energy

Volume

3-D space

Qualitative measurement

describes without using numbers or units. Ex: color, texture, material

Quantitative measurement

describing matter using standardized measurable units. You must have a number and a unit.

Fundamental units

not coming from something else (length, mass, temperature, mole)

Derived units

coming from something else (volume, water displacement, density)

Density of water

1.0g/mL, 1.00 cm^3

Precision

how close your values are to one another when finding an average

Accuracy

how close your value is to the correct value

Significant figures

all digits are significant figures, zeros between the numbers count as S.F, zeros to the right only count if there is a decimal point, and zeros to the left don't count

Calculations using significant figures

when multiplying/dividing, answer must have the least number of S.F using the numbers in the problem, when adding/subtracting, answer must have the least number of digits to the right of the decimal point

Dimensional Analysis

(first number / 1) (unit you are converting to / unit of first number)

Percent Error

(measured value - accepted value / accepted value ) (100) , where the measured value is the experimental value and the accepted value is the known/ published value. Answer must be in absolute value, unless given a (+,-) option

Mole

amount of a substance. There are 6.02 * 10^23 particles of matter in a mole, and the volume of one mole at standard temperature and pressure is 22.4 liters

Gram Formula Mass

found by adding up the atomic mass of each mole of a compound. Units: g/mol

Hydrate

an "ionic" substance that contains water molecules

Stp

standard temperature- 0 C, 273 K
standard pressure- 1 atm, kfa
1 atm \= 760 torr\=101.3 KPA\=760 mmHg

Estp

1 mole of any gas occupies 22.4 L

Diatomic molecules

H2 O2 F2 Br2 I2 N2 Cl2 (HOFBRINCL!)

Mol formulas

mol \= grams(g) / molar mass (g/mol))
mol \= liters(L) / 22.4 (L/mol)
mol \= \# of particles / (6.02*10^23) particles/mol

Matter

anything that has mass and volume. There are two categories of matter: pure substances and mixtures.

Pure substance

always homogeneous. Includes elements, compounds, molecules, and diatomic molecules. Particle diagram is a filled in or empty circle.

Element

a pure substance that can not be broken down by chemical means and contains atoms with the same atomic number and proton number.

Molecule

Electrically neutral group of two or more atoms chemically combined. Atoms can be the same or different. Ex: diatomic molecules, compounds

Diatomic Molecules

(HOFBRINCL) represent by two circles that are the same size and the same color

Compound

two or more different elements/atoms chemically combined, that are electrically neutral. Usually represented by three circles, with one big one in the middle

Mixture

a mixture of two or more substances with different physical and chemical properties that can be separated by physical means. (filtration, distillation, evaporation, chromatography)

Homogeneous mixture

a mixture that does not contain visible particles. It has the same physical and chemical properties throughout. It is a solution (aq)

Heterogeneous mixture

not the same throughout. Ex: oil and water layered

Alloy

homogeneous mixture. 2 different metals, melted, physically combined and cooled. Ex: brass \= copper & zinc, bronze\= copper + tin, steel \= iron + carbon

Phases of matter

solid(s)- different kinds of solids:
crystal (crystalline): regular geometric pattern. Definite volume and shape. *amorphous- ex: glass
precipitate- a solid that when added to a liquid falls to the bottom of the container (beaker). Ex: sand in H2O
liquid(l)- definite volume and indefinite shape (takes the shape of the container)
gas (g)- indefinite volume and indefinite shape (takes the shape of the container for both)

Properties of matter

physical properties- based on observation
Intensive property- does not depend on the amount of matter given.
Ex: color, odor, texture, density, conductivity (the ability to transmit heat and/or electricity), malleability (the ability to be flattened into sheets), solubility (the ability to dissolve), phase of matter, ductility (ability to be drawn/pulled into wire)
Extensive property- does depend on amount of matter given (has a number and unit)
Ex: volume, mass, length/area
chemical properties- how does matter behave in the presence of other chemicals. Ex: flammable, supports combustion, reacts with acid, rusting, corrosion
physical changes- the matter undergoes a physical change, and a new substance is formed. Identity and properties remain the same. Ex: melting, freezing, boiling, dissolving, tearing, crushing
chemical changes- new substance is formed. Ex: combustion, decomposition, rusting
reactants -\> producers, atoms are rearranged

Phase changes

verbs. Fusion/melting (s -\> l), freezing (l -\> s), boiling/evaporating (l -\> g), condensing (g -\> l), sublimation (s -\> g), despotism (g -\> s).

Periodic Table

organized by increasing atomic number/ proton number
Group- vertical column (18 groups)
- elements in the same group have smaller chemical properties based on the same number of valence electrons
Periods-horizontal rows (7 rows)

Group 1 - Alkali metals - 1 valence electron, +1 oxidation number
Do not exist in nature alone
Highly reactive
Group 2 - Alkaline earth metals - 2 valence electrons, +2 oxidation number
Do not exist in nature alone
Reactive
Group 3 to 12 - colored solutions - multiple oxidation numbers
Other metals - between metals and metalloids
Metalloids/ semimetals- properties of metals and/or nonmetals
Group 17 - halogens, halides - 7 valence electrons
Br is cutoff (l), below it is only solids
Phases of matter
Highly reactive
Do not exist alone
Group 18 - noble gasses - 8 valence electrons expect for He2
Most are unreactive gasses (inert)
Row 8 - Lanthanides
Row 9 - Actinides
Atomic number greater than 82 has no stable isotopes - radioactive
Greater than 95 - man-made (nuclear chemistry)

Properties of metals

Have luster, conduct heat/electricity, malleable, ductile, not soluble in water, lose electrons, and oxidation number is positive, solid at room temperature, except for mercury (Hg)

Properties of non-metals

Not malleable, not ductile, does not conduct electricity/heat (except granite), brittle, no luster

Energy

electrical, solar, nuclear, geothermal, chemical, wind, hydro

Kinetic energy

energy is due to the motion of the object

Average kinetic energy

directly proportional to its kelvin temperature
Temperature is a measurement of heat energy, and is not a form of energy

Potential energy

stored energy (in chemical bonds). Heat energy moves/ transfers from high to low

Endothermic

a process where energy is absorbed
Reactant ~\> delta H + ~\> product
Positive- add the delta H to the left
Negative- add the delta H to the right

Exothermic

process where energy is released
Reactant ~\> product + energy

Phase change diagram

1st flat is melting/ freezing, 2nd flat is boiling/ condensing, and phase change occurs on the flats

Dynamic equilibrium

state of balance; constant motion of particles; in transition

Calculating heat energy

q \= mC delta T
Heat \= (mass) (specific heat capacity) (change in temperature)
J \= (g) (4.18 J/g*k) (Tfinal - Tinitial )
4.18 is only for H20

Heat of fusion

amount of energy required to melt/freeze 1 gram of a substance
q\=mHf
heat\= (mass) (heat of fusion)
J \= (grams) ( 334 J/g)
No temp change, it's a phase change

Heat of vaporization

amount of energy required to boil/condense 1 gram of a substance
q\=mHv
Heat \= (mass) (heat of vaporization)
J\= (grams) (2260 J/g)

Separating a mixture (physical)

Filtration - separates solid from a liquid (solid particles need to be large enough, remaining solid is called the precipitate)
Evaporation- used to separate dissolved solids from a liquid
Distillation - used to separate a mixture of liquids with different boiling points
Functional distillation- distillation used to separate components in crude oil
Chromatography- used for separating mixtures. Requires a mobile and stationary phase ( also used to separate substances based on different molarities)

Calormetry

Calorimeter is a device used to measure the heat released or absorbed by chemical reactions
Law of conservation energy- energy is never created nor destroyed, but transferred
Chemical reaction takes place inside the smaller box, while the larger box is filled with water. The temp of the water will change depending on whether the reaction is endo/exothermic
Using q\=mC delta T, we can determine the amount of heat transferred

Standard pressure

1 atm \= 101.3 kPa \= 760 torr \= 760 mmHg

Gas pressure

force on a given area; the force of the gas molecules hitting the wall of the container

Boyle's Law

for a given mass, at a constant temperature, the volume of the gas varies inversely with pressure
P1V1\=P2V2
If the container is made larger and the particles move at the same speed (same temp) they will hit the walls of the container with less force and less often

Charles' Law

the volume of a fixed mass at a constant pressure is directly proportional to its kelvin temperature
If the volume increases, the temperature increases, if the volume decreases the temperature decreases
V1 / T1 \= V2 / T2
If the volume is increasing, the speed (temperature) of the molecules must increase in order to hit the wall of the container at equal force (pressure)

Guylussac's Law

P1 / T1 \= P2 / T2
Directly proportional

Combined Gas Law

(P1)(V1) / T1 \= (P2)(V2) / T2

Ideal Gas Law

no comparison; allows us to evaluate all the variables in one equation
PV \= nRT
P \= pressure (atm)
V \= volume (L)
n \= moles
R \= constant (0.08206 L * atm/mol * k)
T \= temperature (K)

Density of a gas not at STP

d \= PMM / RT , MM \= dRT / P

Density of a gas at STP

d \= g / L , MM / 22.4 L/mol

Dalton's Law of Partial Pressure

for a mixture of gasses in a container the total pressure is the sum of the pressures that each gas would exert if alone
Ptotal \= P1 + P2 + P3 + ...

Kinetic Molecular Theory

ideal gas; a model that attempts to explain the properties and behavior of an ideal gas
1. The particles are so small compared with the distances between them that the volume of the individual particles can be assumed to be negligible (zero)
2. The particles are in constant random straight-line motion. The collision of the particles with the wall of the container is the gas pressure
3. The particles are assumed to exert no force on each other. They are assumed to neither attract nor repel each other.
4. The average kinetic energy of the particles is directly proportional to its kelvin temperature
Ideal conditions:
High temperature
Low pressure
Low molar mass (He, H2)

Real Gas; Kinetic Molecular Theory

1. Gas particles do have volume (finite)
2. Do exert force of attraction on each other especially if the temperature is low

Graham's Law of Effusion / Diffusion

Diffusion- mixing of gasses
Effusion - term used to describe the passage of a gas through a tiny hole
Rate of effusion measures the speed at which a gas transfers through
R1 / R2 \= √MM2 / MM1

Phase Diagram H2O

curve on the H20 phase diagram (s) (l) (g) line slants left due to the density of H2O (s) < H20 (l). All the other diagrams slant right

Vapor Pressure

Vapor - substance that is a gas but at the given temperature should be a liquid or solid
Ex: H20 at 50 C should be a liquid but might have enough energy to escape the surface of the liquid to become a gas
Vapor pressure - pressure exerted by the vapor back down on the liquid or solid
Equilibrium pressure above its liquid or solid
At equilibrium the rate of evaporation \= the rate of condensation
Increasing temperature of a given substance increases the water pressure

Atom

simplest form of matter
Proton- located in the nucleus, +1 charge, 1 a.m.u mass (Table O)
Neutron- located in the nucleus, no charge, 1 a.m.u (slightly larger than proton)
Electron- located outside the nucleus (orbital- region of probable location), -1 charge, 0 or 1/1856 a.m.u mass

Determining \# of protons, neutrons, and electrons

Atomic number \= \# of protons
Identifies the element
\# of protons \= \# of electrons in a neutral atom
Mass \# \= number of protons + number of neutrons
Carbon is standard for table

Isotope

heavy or light form of the same element that occurs due to a different number of neutrons

Hydrogen isotopes

Hydrogen/protium - 1 proton, 0 neutrons, 1 a.m.u mass
Deuterium - 1 proton, 1 neutron, 2 a.m.u mass
Tritium - 1 proton, 2 neutrons, 3 a.m.u mass

Weighted average

the weighted average of the relative abundance of the naturally occurring isotopes
(abundance * atomic mass + abundance * atomic mass) / 100

Ion

charged electron that occurs when atoms gain or lose electrons
Cation - atoms that lose an electron are positive (metals)
Anion - atoms that gain an electron are negative (usually non-metals)

Charge on an ion

\# of protons - \# of electrons

Mordern Orbital Theory

also known as quantum theory
Orbital - region of probable location of the electron
There are 7 (n) principle energy levels (shells)
Each principal energy level has sublevels (s,p,d,f)
Each sublevel contains one or more orbitals that can hold 0, 1, or 2 electrons
When two electrons occupy the same orbital they have an opposite spin (+-) and repel each other
S has 1 orbital, d has 5, p has 6, f has 7

Expanded electron configuration

used because after \#20 electrons do not fill in order

Orbital diagrams

Hund's rule - each orbital in the same sublevel gets one electron before adding a second
Valence electrons - electrons found in the outermost PEL
Max is 8
Shell \= PEL \= same period/row
Lewis dot diagram
S and p electrons only
H + He - duet rule, 2 valence electrons
All elements want to be a noble gas and obtain 8 electrons - octet rule

Isoelectronic

different elements with the same electron configuration
Elements in the same group have similar chemical properties

Ground state vs. excited state

Ground state electrons have the lowest possible energy for an atom/molecule
Periodic table fills in order, one electron for every proton is added - Aufbau principle
When energy is added to a ground state electron it will move/jump to a higher sublevel/PEL. This is the excited state, and it is very unstable
Energy added is usually heat or electricity

Wavelengths / line spectrum

serve to identify the element
Each element has its own unique bright line spectrum (spectral lines) or "fingerprint"
Light is seen when the electron falls back down from excited state to ground state from high energy to low

History of the atom

Dalton - 1800
Matter is composed of indivisible particles called the atom. Atoms are indestructible (untrue - atoms are made of protons, neutrons, and electrons and can be divided)
All atoms of an element are identical in mass and properties (untrue - isotopes have different masses)
Compounds are formed by a combination of two or more atoms
A chemical reaction is a rearrangement of atoms
JJ Thomson/Crooke - early 1900's
Cathode ray tube experiment
Revealed that the electron is negative
Plum pudding model
Claimed that there were positive and negative charges embedded in one sphere
We now know that electrons are outside the nucleus
Rutherford
Small dense positively charged nucleus containing most of the atom's mass
Large volume of empty space around the nucleus in which the electron moves
Gold foil experiment
Bohr
Planetary model
Electrons move around the nucleus in a fixed orbit (wrong - they are orbitals and not fixed) with constant energy
Current
Wave mechanical model
Quantum mechanic model

Ionization energy

the amount of energy required to remove an electron from gaseous atom (endothermic)
Trend - ionization energy decreases as you move down a group
Why? - as you go down a group the outermost electron is further away from the nucleus (proton +, electron -). The outermost electron becomes easier to remove because it is not held as well. The more PELs added the more the nuclear charge is shielded and blocked from the outermost electrons, the easier it is to remove an electron, the lower the ionization energy. The lower the ionization energy the more reactive/active the metal element (metals lose electrons)
Most active metals are found at the bottom of the table
Trend - as you go across a period, ionization energy increases
Why? - as you go across a period, the outermost electron is in the same PEL and is almost the same distance from the nucleus. As you go across a period the number of protons increases, and as the nuclear charge increases it holds onto the electron making it harder to remove

Electronegativity

the ability of an atom to attract the electron of another atom in a molecule
Based on Pauling scale
Trend - across a period the electronegativity increases
Why? - as you go across a period nuclear charge increases. It has a greater ability to attract an electron (elements are in the same period and are at an equal distance from the nucleus)
Trend - down a group electronegativity decreases
Why? - due to the increase in the number of PELs, the nuclear charge is shielded from allowing the electron from the other atom to be attracted
Shielding effect - when the inner PELs block the nuclear charge

Atomic radius

half the distance between the nuclei in a molecule containing identical atoms
Trend - as you go down a group atomic radius increases
Why? - number of PELs increases
Trend - across a period atomic radius decreases
Why? - as the number of protons increases, nuclear charge increases, drawing in the outermost electron, making the radius smaller

Discontinuties

trends aren't always followed
Shielding or blocking an electron from the nuclear charge makes it easier to remove, lowering ionization energy
Electron repulsion in the sublevels causes a lower ionization energy (electrons repel each other, due to both being negative)

Atomic radius: ions

Cation - positive ion
The atomic radius of an atom is larger than its positive ion because it lost a PEL
Anion - negative ion
The atomic radius of a negative ion is larger than the atom. Extra electoral in the orbital causes electron repulsion

Oxidation numbers

the oxidation number of an element is the apparent change assigned to that atom in a molecule

Oxidatiation numbers : Rules

The sum of the oxidation numbers in a molecule is zero (neutral)
The charge on an uncombined free element is zero (all elements in the periodic table are negative)
Special rules:
Hydrogen \= usually +1 when first in a compound, and \=[pppppppp yyyy 2w

Bianary compound

a compound with two different elements

Tertirary compound

a compound with more than two elements and usually contains a polyatomic ion unless it is an organic structure

Polyatomic ion

a group of atoms that are a unit but have a charge

Solving for unknown oxidation numbers

First number is always positive and the second number is always negative, unless there is a polyatomic ion
Exception : if the electronegativity of the first element is higher, than it is negative
The whole compound must equal zero

Element

one type of atom that cannot be broken down in a chemical reaction

Compound

two or more different elements chemically combined

Ionic compound

electrons are transferred from one element to another
A metal loses electrons and reacts with a non-metal, which gains electrons
Conducts electricity when melted and is a solution

Covalent compound

molecular compound that occurs when two different elements, both nonmetals, share electrons
Electrons are not actually transferred (they are in the middle of the two elements), there are no ions
This is an apparent change

Rules for naming compounds

If the first element in the compound has multiple oxidation numbers, solve for the oxidation number and assign it a roman numeral that corresponds to it
Write the name of the second element and change the ending to -ide
Ex: hydride, oxide, bromide, sulfide, nitride, etc

Writing chemical formulas

Find the charge on both elements (remember that roman numeral represents the charge on the first element)
If one of the names is a polyatomic ion put it in parentheses
Cris-cross the charges and drop the sign
Reduce if necessary. Do not reduce inside parentheses

IUPAC prefix method

uses prefixes attached to the second element in the compound to characterize the compound. Based on how many of the second element there are
1 \= mono, 2 \= di, 3 \= tri, 4 \= tetra, 5 \= penta, 6 \= hexa
Ex: CO2 \= carbon dioxide

Converting polyatomic ions to acids

if the polyatomic ion ends in -ate, it becomes an acid that ends in -ic. If it ends in -ite, it becomes an acid that ends in -ous.
Ex: sulfate → sulfuric acid
Sulfite → sulfuros acid

Empirical and molecular formula

Molecular formula - the exact formula of a molecule, giving the types of atom / elements / polyatomic ions and the number of each
Empirical formula - the simplest whole number ratio of the elements

Calculating empirical formula

If given a percentage, assume 100 grams and convert percentage to grams
Divide each substance (element) by its molar mass (moles)
Go out 3 - 4 decimal places for answer, don't round for molar mass
Divide each answer from step 2 by the smallest number in step 2
If all the numbers in step 3 are whole numbers, use them as subscripts. If not, multiply them by a common denominator to make them whole numbers
To calculate the molecular formula:
Molecular mass / empirical mass (mass calculated using empirical formula) \= a whole number to multiply with the subscripts from step 4