Chem Test

Front 1

trends

Back 1

1. size and radius of atoms and ions
2. ionization energy
3. electron affinity
4. electronegativity

Front 2

effective nuclear charge

Back 2

(Zeff) the charge experienced by a particular electron in a many electron atom.

Front 3

calculation of effective nuclear charge

Back 3

Zeff = Z - S
where Z= atomic #/nuclear charge and S=avergae # of screening e-

Front 4

(ex) for Na

Back 4

Z=11 (b/c 11 protons)
S=10 (b/c 10 e- in between)
Z - S = 11 - 10 = +1

Front 5

Trend (1):
size of atoms and ions

Back 5

when atoms interact they:
1. collide and ricochet off (no bond)
2. collide and bond

Front 6

trends for size of atoms

Back 6

1. down a group, atoms become larger: n increases, so distance between nucleus and outermost e- increases so atomic radius and size increases
2. across a period (from L to R) the size decreases b/c the Zeff increases

Front 7

trends for ions

Back 7

1. cations are smaller than their parent atom
2. anions are larger than their parent

Front 8

Trend (2):
Ionization Energy

Back 8

(I) the minimal energy required to remove an e- from a ground state of a gaseous atom or ion.

Front 9

first ionization energy

Back 9

(I1) amount of energy required to remove an e- from a gaseous neutral atom

(ex) Na -- Na+ + e- I1=495 kj/mol (get from table)

Front 10

2nd ionization energy

Back 10

(I2) energy required to remove second e- from an ion not an atom

Front 11

trends for I:

Back 11

1. generally increases across a period
2. generally decreases as we go down a group

Front 12

exceptions to ionization trends

Back 12

1. removal of the 1st and 4th p e-; removal of 1st p e- decreases energy b/c better shielding of nucleus in s orbital b/c nucleus in completely surrounded so for ion (s2p0) gives us lower energy = lower I
2. when 4th e- is added to p and paired an increase in energy b/c they repel each other so the removal of the 4th e- = decrease in energy, lower I.

Front 13

Trend (3):
Electron Affinity

Back 13

energy change that occurs from adding an e- to a gaseous atom

Front 14

trends for electron affinity

Back 14

1. increase from L to R across period
2. does not change much as you go up or down a group

Front 15

greatest affinity =

low affinities =

Back 15

halogens (7A)b/c only need one e- to get stable nobel gas configuration = lower energy

low affinities for 5A due to adding the e- to half filled p shell

Front 16

Trend (4):
Electronegativity

Back 16

measure of the ability of an atom in a bond to attract e- to itself

Front 17

trends for electronegativity

Back 17

1. increases across of period from L to R
2. increases up a group; fluorine is most electronegative element; Cs is least; C and H have almost the same electronegativity

Front 18

group trends

Back 18

1. metals
2. nonmetals
3. metalloids

Front 19

trend for metalic character

Back 19

decreases from L to R

Front 20

Metals

Back 20

shiny, lustrous, various colors, sold metals are malleable (hammered into thin sheets) and ductile (pulled into thin wires); good conductors of head and electricity; most metal oxides form ionic solic which forms basic solution; tend to form cations in solution

Front 21

nonmetals

Back 21

not lustrous, many colors; solids are usually brittle; some hard and some soft; poor conductors; nonmetal oxides are usually molecular substance which form basic solution; tend to form anions in solution

Front 22

metalloids

Back 22

share properties of metals and nonmentals to varying degrees.

Front 23

family trends:
(1) Alkali metals

Back 23

1. group 1A
2. soft w/ low melting points
3. low I1 wh/ allows them to make +1 ions easily
4. very reactive
(a)combine directly with most nonmetals
(b)form metal hydrides, metal sulfides, and metal halides
(c) react vigorously w/ water

Front 24

family trends:
(2) Aldaline Earth Metals

Back 24

1. group 2A
2. low I1 but not low as 1A's; less reactive
3. the heavier the metal the more reactive

Front 25

examples:

Back 25

1. Be: no reaction with H2O liquid or with H2O gas
2. Mg: no reaction with H2O liquid but will react low with H2O gas
3. Ca, Sr, Ba: slow reaction with H2O liquid and reaction with H2O gas

Front 26

family trends:
(3) Hydrogen

Back 26

1. normally seen as H2 gas: under extreme pressure can be made metallic
2. combustable
3. prefers to share e- in cavalent bonds in molecular compounds
4. reacts with metals and nonmetals
(a)nonmetals = exothermic
(b)metals = forms solid metal hydride
(c) will lose e- to form H+ seen mostly in sol

Front 27

family trends:
(4)Oxygen

Back 27

1.nonmetallic to metallic as move down group
2. O2 is typical form (other 6A's are solid)
3. ozone (O3)
4. forms O2- ions
5. forms O2 2- peroxide ions and O 2- superoxide ions

Front 28

sulfur

Back 28

1. most stable = S8 solid

Front 29

family trends:
(5) Halogens

Back 29

1. group 7A
2. nonmetals, diatomic
3. easily gain e- to form -1 ions
4.F2 is most reactive and strongest oxidizer
5. Cl2 is very reactive but less than F2
6. react with most metals to form ionic halides

Front 30

family trends:
(6)Nobel Gases

Back 30

1. group 8A
2. nonmetal
3. highly unreactive; don't readily lose e-; have octet e- configuraton

Front 31

electron configuration for ions

Back 31

1.cation:remove e- from elemental configuration: remove e- from orbital with largest n.
2.anions: add e- to elemental configuration with lowest available n

Front 32

3 types of bonding

Back 32

1. transfer e- form one atom to another: Ionic Bonding
2. sharing e- b/t atoms: Covalent Bonding
3. delocalization of e- over whole solid: Metallic Bonding

Front 33

Lewis symbols

Back 33

way of depicting valence e- for bonding; disperse e- around symbol using hund's rule; for representative elements

Front 34

Bonding Basics

Back 34

1. octet rule: reactivity and bonding generally work toward an octet or nobel gas configuration
2.nobel gases a full outer shell (8e-) except for He which has 2 e- in full shell
3. other atoms tend to gain or lose e- or share e- to reach this stable full outer shell conf.

Front 35

ionic bonding

Back 35

energy released in reaction forming stable product; bond formed by attraction between positive and negative charges

Front 36

lattice energy

Back 36

use to break up; solid network or crystal lattice formed by ions; lattice energies = large positive #'s

Front 37

Eel =

Back 37

KQ1Q2/d where:
k = 8.99 X 10^9 jm/c^2
Q1,Q2 = charge of ions
d = distance (m)

Front 38

why do ions form as they do?

Back 38

Na=[Ne]3s^1
Na+=[Ne] stable nobel conf.
to remove a 2nd e- comes from inner core; requires more energy than lattice energy provides

Front 39

nonmetals...

Back 39

gain of e- is either exothermic or slightly endothermic if added to a valence shell.

Front 40

example

Back 40

Cl=[Ne]3s^23p^5
Cl=[Ne]3s^23p^6 to add 2nd e- it would have to be added to higher energy shell and energy input would exceed lattice energy

Front 41

covalent bonds

Back 41

sharing e-, usually b/t nonmetals; most substances are covalent; nuclei repel each other; e- repel each other, but nuclei attract the e- from other atom; when atoms come close enough (bonding distance) then the opposing nuclei attracts opposing e- and e- are shared

Front 42

polyatomic ions

Back 42

2 or more atoms bound together primarily by covalent interactions to form stable group w/ + and - charges

Front 43

covalent bonds and polarity

Back 43

1. w/ covalent bonds if e- shared equally: non-polar covalent bond
2. covalent bond where e- drawn or attracted to one atom more than another: polar covalent bonds

Front 44

electronegativity scale

Back 44

1. < 0.5: nonpolar covalent
2. b/t >/= 0.5 & < 2.0: polar covalent
3. >/= 2.0: ionic bond

Front 45

polarity

Back 45

1. electronegativity used to measure
2. points to more electronegative atom

Front 46

dipole moment

Back 46

measure of polariy of a molecule
u = Qr where r = bonding radius

Front 47

formal charges

Back 47

help decide b/t reasonable structures

Front 48

calculating formal charges

Back 48

FC=[#valence e- for atom] - [# nonbonding e- on atom] - [1/2 #bonding e- on atom]

Front 49

reasonable structures

Back 49

they are stable as either and interchange (or resonate) b/t the structures
example: ozone

Front 50

exceptions to octet rule

Back 50

1. molecules with odd # of total valence e-
2. molecules which central atom can exist stably w/ less than an octet
3. molecules where central atom can accomidate more that an octet

Front 51

bond energy

Back 51

strength of bonds determined by the energy necessary to break the bond

Front 52

bond enthalpy

Back 52

always requires energy to break bond; always positive

Front 53

bond length

Back 53

the distance b/t the nuclei of bonding atoms

Front 54

bond order

Back 54

a measure of the # of bonding e- pairs b/t atoms

1.single bond= bond order 1
2.double bond= bond order 2
3. triple bond= bond order 3