Chem 7

Front 1

Electronic structure

Back 1

the energy and location of each electron in an atom

Front 2

Chemical properties of elements are related to ...

Back 2

electronic structure

Front 3

Wave length

Back 3

(λ) is the distance between identical points on successive waves

units m, cm, nm

Front 4

Frequency

Back 4

(ν) is the number of waves that pass through a particular point in one second

units - Hz (cycles/s)

Front 5

Amplitude

Back 5

the vertical distance from the midline of a wave to the peak or through

Front 6

Speed =

Back 6

λ ⋅ V

wave length × frequency

Front 7

Electromagnetic radiation

Back 7

(radient energy) emission and transmission of energy in the form of electromagnetic waves

Front 8

Speed of light

Back 8

(c) = 3.00 x 10^8 m/s

speed of light in a vacuum
is a physical constant

Front 9

Quantum

Back 9

the smallest quantity of energy that can be emitted (or absorbed) in the form of electromagnetic radiation

Front 10

h=

Back 10

Planck's constant
= 6.626 x 10^-34 J∙s

Front 11

Energy of a single quantum

Back 11

E=hv
Planck's Constant x frequency

Front 12

The change of energy for a system ∆E can be represented by the equation..

Back 12

∆E=nhv
n= interger (1,2,3,...)
h= Planck's contant
v= frequency of electromagnetic radiation absorbed or released

Front 13

Photon

Back 13

particles of electro magnetic energy with energy E proportional to the observed frequency of light

E=hv=hc/λ

Front 14

Bohr model

Back 14

quantum model for the hydrogen atom. Proposed that the electron in a hydrogen atom moves around the nucleus only in certain allowed circular orbits

Front 15

Equation for energies corresponding to each allowed orbit for the electron in a hydrogen atom

Back 15

E= -2.18x10^18 J (1/n^2)
n= principal quantum number

Front 16

Ground state

Back 16

the lowest energy state
(e- in the n=1 orbit)

Front 17

Excited state

Back 17

higher energy state

Front 18

Equation for a change in energy ∆E when the electron changes energy state

Back 18

∆E= energy of final state - energy of initial state
∆E=
=hv
=h(c/λ)
=-2.18x10^-18J[(1/nf^2) - (1/ni^2)]

Front 19

Wavelength of a emitted photon can be calculated from the equation...

Back 19

∆E=h(c/λ) or λ=hc/∆E

Front 20

De broqlie relation

Back 20

λ=h/m⋄v
m=mass
v=velocity

Front 21

Heisenberg uncertainty principle

Back 21

it is impossible to know simultaneously both the momentum (p=m⋄v) and the position of a particle with certainty

Front 22

Quantum mechanics

Back 22

the branch of physics that mathematically describes the wave properties of submicroscopic particles

Front 23

Ψ

Back 23

wave function

Front 24

Ψ^2 ∝

Back 24

probability of finding the electron in a certain region of space

defines the distribution of electron density in a 3D space around the nucleus

Front 25

Electron density

Back 25

gives the probability that an electron will be found in a a particular region of an atom

Front 26

Atomic orbital

Back 26

3D volume of space where there is a high probability of finding e^-

Front 27

Principle quantum number (n)

Back 27

has integral values (1,2,3...)
related to size and energy of the orbital
As n increases the orbital becomes larger and the elctron spends more time farther from the nucleus.
An increase in n also means higher energy, because the electron is less tightly bound to the nucleus, and the energy is less negative.

Front 28

Orbitals of the same quantum state n are said to belong to same...

Back 28

shell

KLMN
1234

Front 29

Azimuthal quantum number

Back 29

(ℓ) can have any integer value from 0...(n-1) fro each value of n.
N=1 ℓ=0
N=2 ℓ=0,1
N=3 ℓ=0,1,2
N=4 ℓ=0,1,2,3
The shape of the orbital depends on ℓ
For a given n, the energy of the orbital increases with increasing ℓ

Front 30

Orbitals of the same n but different ℓ are said to belong to different ...

Back 30

sub shells of a given shell
ℓ=0 Subshell=s
ℓ=1 Sub shell=p
ℓ=2 Sub shell=d
ℓ=3 Sub shell=f
ℓ=4 Sub shell=g

Front 31

Magnetic quantum number

Back 31

(m sub ℓ) the allowed values are integers from -ℓ to +ℓ
ℓ=0 m= 0
ℓ=1 m= -1,0,1
ℓ=2 m= -2,-1,0,1,2
ℓ=3 m= -3,-2,-1,0,1,2,3
The spacial orientation of orbitals depends on m

Front 32

Shielding effect

Back 32

electrons in the inner shells block electrons in the outer shells from the stabilizing positive charge of the nucleus

Front 33

Penetration effect

Back 33

the greater the probability the electrons can be found close to the nucleus, the more penetrating the orbital
s>p>d>f

Front 34

Closer to the nucleus, ____ attraction, ____ energy.

Back 34

more
less

Front 35

More penetrating the ___ energy

Back 35

lower

Front 36

Spin magnetic quantum number

Back 36

(m_s_) the possible values for the spin quantum number are +½ and -½
The orientation of the spin axis of the electron depends on m_s_

Front 37

Aufbau principle

Back 37

electrons occupy orbitals so as to minimize the energy of the atom

Front 38

Hond's rule

Back 38

when filling a sub shell, one electron is placed in each orbital until all are half filled

Front 39

Pauli exclusion principle

Back 39

electrons must be of opposite spin in order to simultaneously occupy the same orbital

Front 40

Magnetic properties provide direct evidence for electron configuration...

Back 40

Paramagnetic
-attracted by magnet
-contains unpaired electrons
Diamagnetic
-repelled by a magnet
-all electrons are paired

Front 41

Valence electrons

Back 41

electrons in the outermost occupied shell of the atom (the electrons involved in chemical bonding)

Front 42

Valence shell

Back 42

the outermost occupied shell of the atom

Front 43

Core electrons

Back 43

inner electrons that are not involved in chemical bonding

Front 44

Atomic radius

Back 44

factors that determine the size of the outermost orbital in an atom ( and thus the atomic radius)

Front 45

The larger the principle quantum number n the ____ the size of the orbital

Back 45

larger

Front 46

The larger the effective nuclear charge the ____ the size of the orbital

Back 46

smaller

Front 47

Effective nuclear charge =

Back 47

actual nuclear charge - shielding effect

Front 48

Anion is ___ than atom

Back 48

larger

Front 49

Cation is ____ than atom

Back 49

smaller

Front 50

Isoelectronic

Back 50

same number of electrons and similar electron configuration

Front 51

Ionization energy

Back 51

the energy required to remove the outermost electron from an isolated gaseous atom

Front 52

_____ have the lowest ionization energy and ______ have the highest ionization energy.

Back 52

metals
nonmetals

Front 53

Exceptions in increasing ionization energy

Back 53

2a → 3a
5a → 6a

Front 54

Electron affinity

Back 54

the energy change that accompanies the addition of an electron to an isolated gaseous atom

Front 55

From left to right across period electron affinity becomes _____ /_____.

Back 55

larger/ more negative
(metals → nonmetals greater tendency to gain electrons)

Front 56

Nobel gases do not form anions therefore _____ and ______.

Back 56

do not gain electrons and have a positive electron affinity

Front 57

E.A exceptions

Back 57

Be
Mg
N

Front 58

Metal Characteristics

Back 58

-tend to have low ionization energy therefore tend to form positive ions relatively easily
- most metal oxides are ionic solids that are basic
- tend to form cations in aqueous solutions

Front 59

Nonmetal characteristics

Back 59

-no luster
- solid nonmetals are very hard (diamond) or soft powders (phosphorus, sulfur)
- poor conductors of heat and electricity
- tend to have large, negative electron affinities and therefore tend to gain electrons relatively easily
-most nonmetal oxides are acidic