chem exam 1

linear VSEPR structure

2 electron regions, no lone pairs, 180 degress

trigonal planar VSEPR structure

3 electron regions, no lone pairs, 120 degrees

bent VSEPR structure

3 electron regions, 1 lone pair, <120 degrees

tetrahedral VSEPR structure

4 electron regions, no lone pairs, ~109 degrees

triagonal pyramidal VSEPR structure

4 electron regions, 1 lone pair, <109 degrees

bent VSEPR structure pt2

4 electron regions, 2 lone pairs, <<109 degrees

trigonal bipyramidal VSEPR structure

5 electron regions, no lone pairs, 120 degrees, 90 degrees

see-saw VSEPR structure

5 electron regions, 1 lone pair, <90 degrees, <120 degrees

t-shape VSEPR structure

5 electron regions, 2 lone pairs, <90 degrees

linear VSEPR structure pt 2

5 electron regions, 3 lone pairs, 180 degrees

octahedral VSEPR structure

6 electron regions, no lone pairs, 90 degrees

square pyramidal VSEPR structure

6 electron regions, 1 lone pair, <90 degrees

square planar VSEPR structure

6 electron regions, 2 lone pairs, 90 degrees

t-shape VSEPR structure

6 electron regions, 3 lone pairs, <90 degrees

linear VSEPR structure

6 electron regions, 4 lone pairs, 180 degrees

formal charge formula

(valence electrons)-(lone pair electrons)-(number of bonds)

atomic radius trends

• decreases from left to right across period (effective nuclear charge increases and pulls electrons closer to nucleus)

• increases up to down across group (each row adds a new electron shell making it larger)

effective nuclear charge trends

• increases from left to right across period (across a period more protons are added to the nucleus and electrons are added to the outer shell, the outer electrons feel a stronger pull which decreases the atomic radius)

• stays the same down a group

ionization energy trends

the minimum energy needed to remove a electron from a gaseous atom or ion

• increases from left to right across a period (number of protons increases so the nuclear charge increases and pulls the electrons tighter, making them harder to remove)

• decreases down a group (atoms get larger when new electron shells are added which means less energy is required to remove electrons)

metals

conduct heat and electricity well, malleable, ductile, all shiny but mercury

nonmetals

poor conductors of heat and electricity, most gases at room temp, can also be liquids and gases

metalloids

have some properties of metals and nonmetals, tend to have physical properties of metals but w/ chemical properties of nonmetals Boron (B), Silicon (Si), Germanium (Ge), Arsenic (As), Antimony (Sb), Tellurium (Te)

transition metals

elements in groups 3-12

noble gases

elements in group 18 of the periodic table

alkali metals

group 1 elements except hydrogen, when they combine with oxides and react with water, they form alkaline (basic) solutions

halogens

group 17 elements, form 1:1 binary ionic compounds w/ alkali metals

alkaline earth metals

group 2 elements, form ionic compounds w/ halogens

chalcogens

group 16 elements

rutherford experiment

a gold foil surrounded by screens of zinc sulfide is blasted by a alpha particle emitter which emits a ray of alpha particles. a majority of the particles go through the gold foil, which shows that atoms are mostly empty space. some alpha particles deflected against the gold foil, indicating that there is a small dense center (nucleus) in a atom. the significance of this experiment disproves the previous plum pudding model and also demonstrates the existence of a nucleus.

mass number

total number of protons and neutrons in the nucleus (can be found on the periodic table by rounding the molar mass to the nearest whole number)

atomic number

number of protons

isotopes

contain the same number of protons but different number of neutrons

ex. ²⁰Ne (10p, 10n) and ²²Ne (10p, 12n)

celsius to kelvin

add 275.15 to the celsius

kinetic energy

½ (mass (in kg))(velocity (in m/s))²

principal quantum number (n)

indicates the the group of orbitals in an atom

angular momentum quantum number (l)

value ranging from 0 to (n-1), that defines shape of orbital

• 0=s, 1=p, 2=d, 3=f, 4=g

magnetic quantum number, (ml)

value from -l to +l that defines the orientation of an orbital in the space around the nucleus of an atom

spin quantum number (ms)

either -1/2 or +1/2 to indicate spin orientation of an electron, up or down

cathode

negative terminal of high-voltage power supply

anode

positive terminal of high-voltage power supply

mass spectrum

graph of data from a mass spectrometer with peaks whose heights are measures of the number of ions of different masses produced from a sample in the spectrometer’s ion source

photoelectric effect

the release of electrons from a material as a result of electromagnetic radiation striking it

threshold frequency

the minimum frequency of light required to produce the photoelectric effect

electron affinity

energy change that happens when an electron is added to a neutral atom to form a negative ion

• across: becomes more negative across a period as effective nuclear charge increases and pulls electron closer

• down: becomes less negative down a group as atomic radius increases and weakens the effective nuclear charge

polar

is asymmetrical, has central lone pairs, and has different surrounding atoms

nonpolar

is symmetrical, doesn’t have central lone pairs, and has identical surrounding atoms

nonpolar bonds

change in electronegativity is less than 0.4, bond between same-element atoms, or a hydrocarbon bond

balmer series

set of hydrogen spectral lines produced when an electron in a hydrogen atom transitions from or to the second energy level

lyman series

a set of ultraviolet spectral lines emitted or absorbed by a hydrogen atom when an electron transitions between the ground state and higher energy states

chromium electron configuration exception

electron configuration is [Ar] 4s13d5 because 2 half filled orbitals are more stable than 1 full one, and it balances electron-electron repulsions