Chem 1-2

Endothermic

Absorbs heat from surroundings and surroundings get colder (melting/heating)

Exothermic

Releases heat to surroundings and surroundings get warmer (freezing/cooling)

Heat change equation

∆Q\=mC∆T

Propagating uncertainty

1. Uncertainty/mass x 100 = ?% for each value
2. Add percents together. Divide by 100
3. Multiply that by answer to question: uncertainty
4. Final answer: answer +/- uncertainty (step 3)

Specific heat

Amount of heat required to raise the temp of 1 gram of a substance 1 degree C. Objects with higher specific heat take more energy to heat. (C)

Random error

Uncertainty

Systematic error

Errors from poor design or procedure

Precision

How close a group of measurements are to each other

Accuracy

How close a measurement is to the true value

Sig figs

1. When doing x or /, answer can't have more sig figs than original numbers
2. When doing + or -, answer can't have more sig figs after the decimal point than the original numbers

Density equation

mass/volume

Density

Degree of compactness of a substance

Heat

Amount of energy used to cause the temp of an object to increase

Temperature

Average kinetic energy

Heat of fusion

Solid to liquid (335 J per g for water)

Heat of vaporization

Liquid to gas (2240 J per g for water)

Phase change

When matter changes from one state to another

Sublimation

Solid to gas

Deposition

Gas to solid

Melting and boiling points for H2O

0 and 100 degrees C

Phase changes and temp

Heat is changing (absorbing/releasing), but temp is not

STP

1. 273 K/0 degrees C
2. 1 bar/100 kpa

1 mole of gas at STP

22.7 L/dm^3

PTV

P=force

T=average KE

V=Amount of space

PV

Inverse

TV

Direct

PT

Direct

Universal gas constant

R

Moles

\# of particles=6.02 x 10^23 atoms/particles/molecules...

Atomic structure

P=Z, E=P

N=A-Z or N=P-E

Mass # (A)=P+N

Atomic # (Z)=P

Isotopes

Atoms of the same element that have different numbers of neutrons

Ions

Charged atoms

Cation

Positive

Anion

Negative

Nuclear decay

The process when the nucleus of an unstable atom loses energy by emitting radiation

Radioactive

Unstable isotopes

Fusion

When atoms join together (2 to 1)

Fission

When an atom breaks apart (1 to 2)

Half life

Time required to reduce something to half of its initial value

Alpha decay

Least harmful

Penetrates skin and paper

4/2a

Most ionizing and least penetrating

Beta decay

Penetrates aluminum

\+=0/+1e

\-=0/-1e

Gamma decay

Most harmful

Penetrates lead and steel

0/0y

Most penetrating and least ionizing

Band of stability

Alpha needs to loose P and N

Beta (+) needs to loose P

Beta (-) needs to loose N

Red and IR

Lowest energy and longest wavelength

Violet and UV

Highest energy and shortest wavelength

Energy levels

Higher levels have more energy, shorter wavelength, and are farther away from nucleus

Ground state

Lowest possible energy level (stable)

Excited state

When an atom absorbs energy, its electrons move to a higher energy level (unstable)

Moving closer to nucleus

Releasing energy

Moving farther from nucleus

Absorbing energy

Electron configuration order

spdf

s is closest to nucleus and lowest energy

Valence electrons

\# of electrons on outer energy level

Period

Left and right

Group

Up and down

How to check electron configuration

\# exponents=# of E

\# exponents in highest energy level=# valence E

Periodic table key

Ex: 3s^2

3=period

s=block

2=number to the right

Electron configuration boxes

1. No more than 2 E and opposite spins in each orbital
2. Put E in each box before pairing up singles
3. Boxes (orbitals) want to be either full or half full, so may need to rearrange

Abbreviated electron configuration

\[Last nobel gas (row farthest right)] then continue with orbitals

How many moles are in \# grams of element?

Mass (g)/molar mass (element)

How many grams are in \# molecules of element?

(Molecules) x (1/6.022 x 10^23) x (molar mass)

How many molecules are in \# grams of element?

(mass (g)) x (1/molar mass) x (6.022 x 10^23)

How many grams are in \# moles of element?

mole x molar mass

First ionization energy

The energy required to remove 1 outermost electron from 1 atom is a gaseous state

Second ionization energy

The energy required to remove the second electron

Periodic properties

Ionization energy, electronegativity, electron affinity, and atomic radius

What do periodic properties depend on?

Electron configuration (P to E ratio) and how strongly Es are attracted to the nucleus

Coulomb's Law

Force and periodic property depends on the distance between the objects (how strongly attracted) and the magnitude of the charges (electron configuration)

Charge and force

Proportional

Distance (r) and force

Inverse

Does distance or charge affect the force more?

Distance

Effective nuclear charge

Net electrical charge acting on an electron.

Bigger # of P increases nuclear charge

Bigger distance decreases nuclear charge because of shielding and it is farther from the nucleus

Shielding

Repulsion of inner electrons acting on other electrons, reducing the effect of the nuclear charge on electrons

Effective nuclear charge equation

Zeff=Z-S

Z=# of P

S=# of non-valence E

Electron shielding effect

Electrons between the nucleus and the valence electrons repel each other, making the atom larger

Electronegatiivity

A measure of attraction one atom has for the electrons of another atom

Atomic radius

The distance from the nucleus to the valence electrons

Electron affinity

The amount of energy released when a neutral gaseous atom gains an electron

Periodic trends

IE, En, and EA are highest at the top right of the table

AR is highest at the bottom left of the table

Ionic radius for cations and anions

Cations=smaller (bc looses a shell)

Anions=bigger (bc gains a shell)

Molecules

Groups of atoms that are chemically bonded together

Diatomic molecules

Two atoms of the same element bonded together (H2...)

H, N, F, O, I, Cl, Br

Have no fear of ice cold beer

Magic 7 bc in 7 shape on table

Molecular compounds

Two or more different elements chemically bonded

Law of constant composition

A compound that always has the same proportions by mass

Law of multiple proportions

A compound that can combine in different ratios to make more that 1 compound

Molecular formula

Indicates the actual number of atoms in a molecule

Empirical formula

Gives the relative number of atoms (ratio)

Structural formula

Shows which atoms are attached to which

Ionic

Metal and nonmetal

Metallic

Metal and metal

Covalent

Nonmetal and nonmetal

Polyatomic ions

Atoms joined in a molecule that together carry a charge

Naming ionic compounds

Metal atom then nonmetal atom ending in ide

Naming covalent compounds

Least EN Non-Metal then more EN Non-Metal ending in ide and Greek Prefixes go in front

Greek prefixes

Mono, di, tri, tetra, penta, hexa, hepta, octa, nona, deca

Naming acids

Hydrogen with covalent compound(s)

Ate or other non polyatomic ions=ic (HCl or HNO3)

Ite=ous (HNO2)

Percent composition

The percent by mass of each element in a compound

Structure of ionic compounds

Ionic crystal lattice

Characteristics of ionic compounds

\-Usually solid

\-Higher melting and boiling points than covalent

\-Conduct electricity in solution state, but not solid state

\-Polar

\-Transfer electrons

Characteristics of covalent compounds

\-Low melting points

\-Do not conduct electricity

\-Share electrons

\-Lack electrons to fill valence shell

Characteristics of metallic compounds

\-Positive ions surrounded by a "sea" or "cloud" of mobile electrons

\-Good conductors of heat and electricity

\-Great strength

\-Malleable and ductile