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Chem/Phys Exam 1 > Chem Review > Flashcards

Chem Review Flashcards

1
Q

Elemental Composition of Human Body

A 
Hydrogen= 63%
Oxygen= 24.2%
Carbon= 10.5%
Nitrogen=1.35%
Major Minerals=0.7%
Trace Elements= 0.01%
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2
Q

Dalton’s Theory of the Atom

A

All matter is composed of atoms that can be neither created nor destroyed.

Atoms of a particular element are identical in size, shape, mass and all other properties and differ from atoms of other elements in these properties.

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3
Q

Atom

A

Element in its smallest form

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4
Q

Molecule

A

Two or more atoms combined (bound)

The atoms may be same element or different elements

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5
Q

Atomic weight

A

Similar to mass

Relative masses of atoms are related to the most abundant carbon isotope, Carbon 12, and differ widely in masses.

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6
Q

Why is the weight of carbon listed as slightly higher than 12 sometimes?

A

Because of the existence of some carbon isotopes in nature, Carbon 13 for example

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7
Q

Neutron

A 
No charge (0)
Amu= 1.0087
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8
Q

Proton

A

Positive Charge +

Amu=1.0078

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9
Q

Electrons

A

Negative charge -

Amu= .00055

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10
Q

Most distinguishing characteristic of an element?

A

The number of protons, or positively charged particles in the nucleus of an atom.
If you lose a proton, you will have a different element

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11
Q

Mass of an element is composed of?

A

Protons and neutrons

Electrons are very small are hardly contribute to the mass (atomic weight)

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12
Q

Atomic Nucleus

A

Atoms have a small concentrated positive charge in the atomic nucleus

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13
Q

What occupies the surrounding space around the nuclear atom?

A

Most of the atom is space and electrons occupy this surrounding space.
Nucleus has a positive charge and electrons has a negative charge and these will be attracted to each other

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14
Q

How much space do electrons and protons take up in atom?

A

Nucleus takes up a very small amount of the atom=protons

Electrons circulation around the nucleus take up the most space

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15
Q

The Neutral Atom

A

Number of electrons=number of protons

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16
Q

Z number of an element

A

Represents Atomic number
Number of protons
If element is neutral, then protons and electrons are equal

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17
Q

How are atoms presented on the periodic table?

A

In their neutral form, although they are not neutral in the environment. This is just an easier way to study them

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18
Q

Can two elements ever have the same number of protons?

A

NO, that’s the distinguishing factor

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19
Q

In the periodic table how are elements arranged?

A

According to atomic number

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20
Q

General notation of elements. What does A, E and Z represent?

A

A= mass (protons+neutrons)
E=element by identity symbol
Z=atomic number (number of protons)

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21
Q

As we move down the periodic table what do we see?

As we move across what do we see?

A

We see similar properties of elements in columns as we move down
We see different properties as we move across?

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22
Q

Isotopes

A

Some atomic masses differ for same atom
Isotopes that have the same atomic number, but different atomic masses.
In isotopes, the atomic nucleus of a given element has a fixed number of protons, but the neutrons vary
Not all isotopes are radioactive

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23
Q

Hydrogen vs. Deuterium (2H)?

A

Hydrogen has 1 proton and 0 neutrons= atomic mass of 1

Deuterium has 1 proton and 1 neutron= atomic mass of 2

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24
Q

Inert or noble gases properties

A

VIIIA (far right of periodic table)
Helium, Argon, Krypton, Xenon, Radon
Properties: do not react, all gases at normal temperature

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25
Q

Alkali or light metals

A

IA, except H- Far left of table
Lithium, Sodium, Potassium, Rubidium, Cesium, Francium
Properties: in pure form have shiny lust, good conduction of heat and electricity

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26
Q

Halogens

A

Group VII A
Fluorine, Chlorine, Bromine, Idoine
Properties: no metals, poor conductors of heat and electricity, all react with hydrogen to form hydrogen halide. Ex: hydrogen fluoride, hydrogen chloride

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27
Q

Transition Metals

A

IIIB-IIB can have variable charges
Row 3 of table- Magnesium, Sulfur.
A decrease in metallic properties is seen.

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28
Q

Atomic Radii of elements

A

Increase in atomic radii from top to bottom of periodic table

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29
Q

The effect of ionization in atomic radii

A

The addition of electrons has a major effect on atomic size

The addition or removal of a single electron roughly increases or decrease the radii by half.

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30
Q

Ionized

A

Gaining or losing an electron

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31
Q

Electronegativity

A

The ability of an atom to draw (attract) electrons or a negative charge towards it.
Electronegativity is not ionization.
Values are relative, note that fluorine has the highest value of 4.0
Electronegativity increases left to right, and decreases from top to bottom.

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32
Q

Why do elements at the bottom of the periodic table has decreased Electronegativity?

A

When a neutral atom has more protons, it has more electrons.
Electrons shield that nucleus and have less interaction with the protons.
More electrons=more shielding=less Electronegativity

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33
Q

Ionization

A

Energy required to ionize an element
Inert elements have the highest ionization energies– most difficult to ionize
Easiest to ionize= Li, Na, K, Rb, Cs

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34
Q

Nuclear Stability

A

Nucleus is packed tightly with positive particles and would seem to want to fly apart
Neutrals act to stabilize the nucleus
In order to achieve stability, the more protons in a nucleus- there needs to be more neutrons.

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35
Q

Ionization energy is a periodic property of elements
Which elements hard to ionize?
Which elements easier to ionize?

A

Hard to ionize or require the most energy to ionize is Helium.
Inert elements are hard to ionize, they have 8 electrons in outer shell and do not like to lose this electron
Easiest to ionize= Alkali metals, Lithium, Sodium, Potassium
Usually give up on electron and become positively charged

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36
Q

When the neutron-to-proton ratio is too high or too low what happens?

A

Too high or too low= nuclear stability is altered leading to radioactive processes.

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37
Q

Radioactive processes

A

Occurs when nuclei are trying to achieve stability
This process changes the number of neutrons or protons to try to reach stability.
When this occurs they release energy in a form of radioactive particles

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38
Q

Beta particle

A

An electron, or -1 charge that is released during radioactive decay process

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39
Q

Alpha particle

A

Associated w/ heavier elements normally

Alpha particle= two protons and two neutrons (+2 charge)

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40
Q

Gamma rays

A

(Y)= no mass or charge, but high energy particles (photons)

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41
Q

When radioactive decay occurs does the element change?

A

YES

Sometimes number of protons is changed. Example: tritium in notes

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42
Q

As elements get larger, what happens to neutron-to-proton ratio?

A

As element gets larger, they are less stable because there are more protons, and nucleus gets bigger. It then is harder to hold this element together.
Neutrons are thought to play a role in stabilizing the protons and keeping them together

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43
Q

Carbon14

A

Is an isotope that is used for carbon dating
1/2 life of carbon 14=5670 years
Half life is the time it takes for radioactivity intensity to decrease by half. You can use this number of determine the age of organic materials
Less carbon 14=older the element is

44
Q

Gamma rays penetration

A

Gamma rays completely penetrate through the body, most worrisome than alpha and beta.
Any exposure to alpha, beta or gamma is related to HOW MUCH you are exposed to

45
Q

What determines an electrons reactivity?

A

The electrons in atoms

46
Q

Neil’s Bohr

A

Proposed that only specific energy levels are possible for electrons in atoms
-electrons circle the nucleus only in certain allowed orbits of specific radii

47
Q

Quantum Mechanics

A

Fixed amount- related to how many electrons can fit in an orbit

48
Q

The energy levels are identified by number. Electron population is determined by what equation?

A

2n(2)

N=energy level, the principal quantum number

49
Q

Energy level of electrons in an atom

A

As N increases, the electrons spend more time farther from the nucleus and are less tightly bound to the nucleus.
Lower energy levels- more stable the atom is

50
Q

In the first energy level, the electrons nearest the nucleus have the highest or lowest energy level?

A

These electrons have the lowest energy level, but they are tightly attracted*

51
Q

As we move up energy level, i.e. From n=1 to n=2, electrons what happens to energy?

A

As you move up energy levels- you are farther from nucleus, these electrons are less attracted to the nucleus but have more energy. These are less stable electrons.

52
Q 
Maximum electron population of each energy level
N=1
N=2
N=3
N=4
A

N=1, 2(1) squared= 2 electrons
N=2, 2(2)squared=8 electrons
N=3 2(3)squared=18 electrons
N=4 2(4)squared=32 electrons

53
Q

Electron Shells (energy levels)

A

Shells or energy levels are composed of subshells, which are composed of orbital.
Shells(energy level or orbitals)>subshells>orbits
Start at:
1(K)
2(L)
3(M)
4(N)

54
Q

Which electrons confer reactivity?

A

KEY POINT: electrons in the outer shell confer reactivity

55
Q

Stability of an element is most closely related to what number?

A

8 is the magic number for stability.
Inert gases have 8 electrons in their outer shell=stable
Most elements want 8 in their outer shell to be stable.

56
Q

Shells- subshells- orbitals

A

Energy level or shells= 1 (K), 2 (L), 3 (M), 4 (N)
Subshells= s, p, d, f
Orbitals= 1, 2, 3

57
Q

S subshell (number of electrons and orbitals)

A

S subshell= 2 electrons total, spherical shaped orbital

one S orbital per energy level

58
Q

P Subshell (number of electrons and number of orbitals)

A

P subshell= 6 electrons, 3 p orbitals per energy level (each orbital can have a maximum of 2 electrons, for a total of 6)
Orbitals exist along different axes: X, y, z
Dumbbell shape

59
Q

D subshell (number of electrons and orbitals)

A

D subshell=10 electrons, five D orbitals exist in energy levels above N=2. (Each orbital can hold a max of 2 electrons for a total of 10 electrons in the d subshell)
Four leaf clover shape

60
Q

F subshell (number of electrons and orbitals)

A

F subshell=14 electrons, 7 orbitals above n=3. Each orbital can hold a max of 2 electrons for a total of 14 electrons.
Complicated shape.

61
Q

Electron probability distribution

A

It is not ever possible to determine exactly where electrons are around the nucleus. Electrons in orbitals are not defined locations, but are distributed based on probabilities of location.
As you move away from the nucleus, electron probably is much lower

62
Q

1s electron density distribution

A

Electron density distribution, such as Hydrogen atom, in the ground state makes a spherical shape of distribution.

63
Q

2p electron density distribution

A

Electron density distribution is a dumbbell shape. 2 electrons are circulating in a figure 8 manner

64
Q

Two electrons in an orbital have similar or opposite signs?

A

Two electrons in an orbital have opposite spins or opposite signs because they want to pair up

65
Q

What happens to the filling of shells as you move up higher energy levels?

A

Sometimes the filling of shells is NOT sequential, meaning you can fill the 4th energy level before we fill the 3rd energy level.

66
Q

Why does the atomic radii decrease in the center part of periodic table, i.e. Transition metals?

A

The electrons in the transition metals live more happily close together, making the atomic radii smaller.

67
Q

Reactivity of electrons in atoms

A

Electrons in the lowest energy level require the most energy to be removed from the atom.
Electrons in the higher energy level require the LEAST energy to be removed from the atom

68
Q

What atoms are most likely going to be involved in chemical reactions?

A

The highest energy level electrons (those in the outer shell).
Not as tightly attracted to nucleus because they’re farther away from the + charge

69
Q

Valence Electrons- Lewis Diagram

A

The dots represent the electrons in the outer shells

These outer electrons are responsible for the valence, or combing capacity of atoms

70
Q

Valence Electrons

A

Electrons that belong to a a valence shell

71
Q

Valence Shell

A

The outermost shell of an atom

72
Q

Valence

A

Number of bonds an atom can form

73
Q

Saturation of Valence

A

There is a limit to combining ability of one atom to another
Related to the number of electrons that are in the outer shell that can shared or bound
Example: Carbon can form only 4 bonds- no more no less, trying to achieve 8 electrons in outer shell for stability

74
Q

Rule of Eight, or Octet Rule

A

Bonds are formed such that each atom is surrounded by a complete Octet (8) of electrons (except hydrogen)
The molecule is more stable with 8 electrons in the outer shell of its atom whether it shares them or not

75
Q

Orbital Notations- Hydrogen

A

One electron in outer shell=

1s(1)

76
Q

Orbital notation for carbon

A

Carbon= 1s(2) 2s(2) 2p(2)
OR
Promotion allows four bonds to form
1s(2) 2s(1) 2px(1) 2py(1) 2pz(1) (shows that there is 4 bonds available for sharing

77
Q

Orbital Notation for Nitrogen

A

1s(2) 2s(2) 2p(3)
OR
1s(2) 2s(2) 2py(1) 2px(1) 2pz(1)

78
Q

Hydroxyl Group

A

O-H

79
Q

Carbonyl or Keto Group

A

Carbon double bonded to oxygen=4 bonds are shared

80
Q

Ene suffix

A

Aminated alkene

Double bond

81
Q

Amine

A

N-H

82
Q

Atoms are more stable when their outer shells are filled… They can do this in two ways?

A
  1. Ionization

2. Electron Sharing

83
Q

Ionization or sharing of elements leads to ______________

A

Bonding with other atoms to form molecules

84
Q

Ionic Bonds

A

Atoms can fill their outer shells by completely acquiring or losing electrons and in doing so become ionized
The consequent electrostatic (+ attraction to a -)interactions result in ionic bonding
Valence electrons are NOT shared

85
Q

Example of sodium and chloride as an ionic bond

A

(Na) will lose an outer electron to form a cation {a cation is positively charged atom or molecule)
(Cl– which is a halogen) will gain an electron to form an anion {an anion is negatively charged atom or molecule}
Idea= atoms w/ opposite charge have the largest attraction

86
Q

A few examples of ionic compounds

A

LiF- lithium fluoride
NaF- sodium fluoride
NaBr sodium bromide
NaI- sodium iodide

  • These are strongly attracted to each
  • Note that these ionic compounds (salts) contain the IA alkali metals and the halogens (7A)
  • These are strongly bound atoms requiring high energy to separate into free ions. They dissolve in water
87
Q

What exactly does stability represent?

A

When atoms achieve stability they have 8 electrons in their outer shell, meaning they can exist and have a longer lifetime

88
Q

Covalent Bond

A

In covalent bonds, electrons are shared between atoms

89
Q

Example of covalent bond: H2

A
  • The sharing of two electrons (a pair) between atoms results in a SINGLE bond or sigma bond.
  • Single bond between two atoms is sausage shaped
  • we generally don’t find pure hydrogen in nature very much
90
Q

Electronegativity

A

Ability of an electron to attract electrons to it

91
Q

Covalent bond example: Chlorine

A

Neutral chlorine has 7 electrons in its outer shell, i.e. 7 valence electrons
– it has single vacancy in the 3p i.e. 1s(2) 2s(2) 2p(6) 3s(2) 3p(5)
orbital that can be filled by sharing another single electron (a single bond)
- chloride gas is relatively stable but not too stable

92
Q

Example of covalent bond: Carbon and Hydrogen

A

Carbon has four valence electrons
Hydrogen has 1 valence electron
Carbon possesses an octet when bound to four other atoms (covalent bond)

93
Q

Examples of common covalent bonds

A 
C-C
O-H
C-N
C-O
Cl-Cl
  • These can only contain sigma bonds, but those capable may contain multiple bonds.
  • Bonds drawn as a single line represent the sharing of two electrons
94
Q

Covalent and Ionic Bonding: Nitrogen and Hydron

A

Nitrogen has 5 valence electrons
When bound to 3 hydrogens it has an octet
However one pair of its electrons is unbound in the octet
The unbound pair of electrons can associate (bind) with a cation such as a proton and the molecule becomes uncharged.

95
Q

Ionic Bonds:

A

Ionic bonds- electrons are transferred
NaCl
Do NOT share electrons

96
Q

Polar Covalent Bonds

A

Electrons are unequally shared

HCl

97
Q

Covalent Bonds

A

Electrons are shared equally

H-H

98
Q

Dipole

A

A molecule with on end having a slightly (partial) negative charge and the other end having a slight positive charge (NOT ionic bonds)

99
Q

Atoms can share more than one pair of electrons resulting in more than one bond between atoms….
Single bond vs Double bond vs Triple bond

A

Single bond= one sigma bond- saturated compound
Double bond= one sigma bond + 1 pi (p) bond
Triple bond= one sigma bond + 2 pi bonds

100
Q

Ethane

A

Single bond

101
Q

EthEne or ethylene

A

Double Bond

102
Q

Ethane (or acetylene)

A

Triple bond

103
Q

Aromatics

A

Alternating double bonds in a ring system

104
Q

When pi electrons (from overlapping p orbitals) delocalize around the ring and make the ring _________

A

More stable, i.e. The pie electrons are in a lower energy state

105
Q

Molecular Nitrogen Formula

A

N2- in atmosphere inert gas common, abundant gas

Molecular nitrogen shares 3 pairs of electrons, therefore it contains a triple bond. One s bond and 2 p bonds

106
Q

Nitrogen Facts

A

Odorless gas, does not support combustion
melting point= -210 degrees C
Boiling point= 196 C

107
Q

Breakdown of the Octet Rule (NO)

A

Nitrogen Oxide (Nitric Oxide)- neutral, but reactive, used as a vasodilator gas

Chem/Phys Exam 1 (1 decks)

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