Chemistry Fundamentals Flashcards
1
Q
Sulfate
A
SO4 2-
2
Q
Sulfite
A
SO3 2-
3
Q
Nitrate
A
NO3-
4
Q
Nitrite
A
NO2-
5
Q
Phosphate
A
PO43-
6
Q
Carbonate
A
CO32-
7
Q
Bicarbonate
A
HCO3-
8
Q
Hydroxide
A
OH-
9
Q
Permanganate
A
MnO4-
10
Q
Chromate
A
CrO42-
11
Q
Dichromate
A
Cr2O72-
12
Q
Ammonium
A
NH4+
13
Q
Cyanide
A
CN-
14
Q
Hyperchlorite
A
ClO-
15
Q
Chlorite
A
ClO2-
16
Q
Chlorate
A
ClO3-
17
Q
Perchlorate
A
ClO4-
18
Q
Order of Magnitude
Ex. By how many orders of magnitude is a centimeter longer than an angstrom
A
An order of magnitude is a factor of 10
1 cm = 10 -2 m and 1 A = 10 -10 m
A centimeter is 8 factors of 10, or 8 orders of magnitude, greater than an angstrom.
19
Q
Density
A
P = mass / volume = m / v
20
Q
Molecule
A
When two or more atoms form a covalent bond, they create a MOLECULE
21
Q
Molecular Formula
A
A compounds molecular formula gives the identities and numbers of the atoms in the molecule. For example, the formula C4H4N2 tells us that this molecule contains 4 carbon atoms, 4 hydrogen atoms and 2 nitrogen atoms.
22
Q
Emperical Formula
A
The REDUCED formula. Ex. C4H4N2 the smallest whole numbers that give the same ration of atoms are 2:2:1. Using these numbers for the atoms, we get the EMPERICAL FORMULA.
23
Q
Formula Weight
A
The sum of the atomic weights of all the atoms in the molecule.
24
Q
Atomic Mass Unit
A
The unit for atomic weight
25
Mole
simply a particular number of things
26
Avogadro's number
27
Conversion from moles to grams
Grams DIVIDED by molecular weight = moles
Moles TIMES molecular weight = moles
28
moles formula
moles = mass in grams / molecular weight
29
Molarity formula
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30
Mole fraction
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31
Limiting Reagent
substance we run out of first
32
Catalyst
a substance that increases the rate of a reaction without being consumed
33
Oxidation State (Oxidation number)
meant to indicate how the atom's ownership of its valence electron changes when it forms a compound
34
Rules for assigning oxidation states
1. The oxidation state of any element in its standard state is zero
2. The sum of the oxidation states of the atoms in a neutral molecule must always be 0 and the sum of the oxidation states of the atoms in an ion must always be equal to the ion's charge.
3. Group 1 metals have a +1 oxidation state, and Group 2 metals have a +2 oxidation state
4. Flourine has a -1 oxidation state
5. Hydrogen has a +1 oxidation state when bonded to something more electronegative than carbon, a -1 oxidation sate when bonded to an atom less electronegative than carbon, and a 0 ox state when bonded to carbon
6. Oxygen has a -2 oxidation state
7. The rest of the halogens have a - 1 oxidation state, and the atoms of the oxygen family have a -2 oxidation state.
35
The order of electronegativities
can be remembered with the neumonic FONClBrISCH
(fawn-cull-brish)
36
Atom
the smallest unit is one ATOM of the element
37
Nucleus
all atoms have a central NUCLEUS, which containes PROTONS and NEUTRONS known collectively as NUCLEONS. Each **proton** has an electric charge of **+1** elem. unit, **neutrons** have **no charge (0)**. NOTE: Nucleus is positive due to this. Outside the nucleus, an atoms contains electrons, and each **electron** has a charge of **-1** elem. unit
38
Atomic Number
The number of protons in the nucleus of an atom is called the **ATOMIC NUMBER, Z** The atomic number of an atom uniquely determines what element the atom is, and Z may be shown explicticly by a subscript before the symbol of the element.
Z A
39
Mass Number
The number of protons plus the number of neutrons in the nucleus of an atom gives the atom's **MASS NUMBER, A.**
If we let N stand for the number of neutrons, then A = Z + N
Mass # = protons plus neutrons
Mass number usually written like A E
40
Isotopes
If 2 atoms of the same elemets differ in their number of neutrons, then they are called **ISOTOPES**
## Footnote
Examples:
7Be (4 protons and 3 neutons)
9Be (4 protons and 5 neutons)
41
Atomic Weight of an element
weighted average of the masses of its naturally occuring isotopes
42
Ion
when a neutral atom gains or loses electrons, it becomes charged, and the resulting atom is called an **ION**
For each electron it gains, an atom acuires a charge of -1 unit, and for each electron it loses it acquires a charge of +1 unit
43
Anion
A **negatively** charged ion
44
Cation
a **POSITIVELY** charged ion
45
Strong Nuclear Force
The protons and neutrons are held together by a force called the **STRONG NUCLEAR FORCE.** Stronger than the electrical force between charged particles since for all atoms besides hydrogen, the strong nuclear force must overcome the electrical repulsion between protons. The strong nuclear force is the most powerful even though it only works over extremely short distances.
\*binds protons and neutrons
46
Radioactive Decay
Unstable nuclei are said to be radioactive and they undergo a transformation to make them more stable. altering the number and ratio of protons and neutrons or just lowering their energy.
47
Weak Nuclear Force
responsible for nuclear decay
48
What are the 3 types of radioactive decay?
Alpha, Beta, and Gamma
49
The nucleus that undergoes the radioactive decay is known as the \_\_\_\_\_\_\_\_, and the resulting more stable nucleus is known as the \_\_\_\_\_\_\_
parent, daughter
50
Alpha Decay
When a large nucleus wants to become more stable by **reducing the number of protons and neutrons in large nucleus**, it emits an alpha particle denoted by 4/2 α, consisting of 2 protons and 2 neutrons. Equivalent to 4/2 He nucleus. Alpha decay **subtracts the parents atomic # by 2** and **subtracts the mass # by 4.**
51
What are the 3 types of Beta Decay?
The 3 types of Beta decay are β-, β+, and electron capture. Each types of beta decay involves the **conversion of a neutron into a proton** through the action of the weak nucear force. NOTE: β particles are worse than α particiles since there less massive. They have more energy and greater penetrating ability.
52
β- decay
When an unstable nucleus contains too many neutrons, it may convert a neutron into a proton and electron.
\* decreases the # of neutrons
\*increases the # of protons
\*Adds 1 to the atomic #
53
β+ decay
When an unstable nucleus containes too few neutrons, it converts a proton into a neutron and positron which is ejected.
54
Electron Capture
An unstable nucleus increases its number of neutrons is to capture an electron from the closest shell and use it in the conversion of a proton into a neutron
\*increases the # of neutrons
\*decreases the # of protons
\*Subtracts 1 from atomic #
55
Gamma Decay
A nucleus in an excited energy state (usualy occurs after a nucleus has undegone alpha or any type of beta decay) can relax toits ground state by emitting energy in the form of one or more photons of electromagnetitic radiation. Gamma photons have no mass or charge and penetrates matter most effectively.
\*Brings an excited nucleus to a lower energy state
\*Does not change mass # or atomic #
56
Half Life
which is denoted by t 1/2 of a radioctive substance is the time it takes for 1/2 of a substance to decay
57
Emission Spectrum
light seperated through its component wavelength. It gives an energentic fingerprint of that element because it consists of a unique sequence of BRIGHT lines that correspond to specific wavelengths and energies.
58
The energies of the photons, or particles of light that are emitted are related to frequency and wavelength by the equation
Energy is DIRECTLY proportional to frequency and INDIRECTLY proportional to wavelenght.
Ex. If energy increases, frequency increases and wavelength decreases
59
Bohr Model of the Atom
Bohr proposed his **quantized** shell model of the atom to explain how electrons can have stable orbits around the nucleus. Bohr modified the Rutherford model by requiring that the electrons move in orbits of fixed size and energy. The energy of an electron depends on the size of the orbit and is lower for smaller orbits. Radiation can occur only when the electron jumps from one orbit to another. The atom will be completely stable in the state with the smallest orbit, since there is no orbit of lower energy into which the electron can jump.
\* Ceiling is 0, closer to nucleus you are getting more and more negative (lowest energy is closest to nucleus)\*
60
Bohr Model Energy States
An electron is initiallly in it ground state (n=1) or its lowest possible energy level. When this electron absorbs a photon it jumps to a higher energy level known as an excited state (n=3). Electrons can relax to ground state immediately or in small jumps.
NOTE: energy absorbed is away from the nucleus, energy emitted brings it closer to the nucleus.
61
Energies of discrete energy levels were given by Bohr in the following equation
We calculate the energy differences between discrete levels by subtracting the initial energy of the electron from the final energy of the electron. We can find the energies of the 2 possible emitted phtons shown above as follows:
62
What is a Bohr atom?
An atom that only contains 1 electron.
63
The Bohr model cannot describe
the electron-electron interactions that exist in many electron atoms
64
How does energy shell change relative to electron energy?
An electron in a higher shell has a greater amount of energy and a greater average distance from the nucleus.
Ex. An electron in the 3rd shell (n = 3) has a higher energy than an electron in the 2nd shell (n = 2) which has more energy than an electron in the 1st shell (n = 1)
65
The energy subshell
In a quantum model of the atom, we no longer describe the path of the electrons around the nucleus as circular orbits but focus on the the probability of finding an electron somewhere in the atom. An orbital describes a 3D region around the nucleus in which electron is most likely to be found.
A subshell in an atom is comprimised of 1 or more orbitals and is denoted by a letter (s, p, d, or f) that describes the shape and energy of the orbitals. Orbitals get more comples and higher in energy in this order. Each energy shell has 1 or more subshells, and each higher energy shell contains 1 additional subshell.
66
The Orbital Orientation
Each subshell contains 1 or more orbitals of the same energy (degenerate orbitals) and these orbitals have different 3D orientations in space. The # of orientation increases by 2 in each successive subshell. For example, the ***s*** subshell contains 1 orientation and the ***p*** subshell contains 3 orientations
67
Recognize the shapes of the orbitals in the s and p subshells
Each s subshell has just 1 spherically symetrical orbital. Each p subshell has 3 orbitals, each depicted a a dumbell with different spatial orientation.
68
Electron Spin
Every electron has 2 possible spin states which can be considered the elctrons intrinsic magnetism. Because of this every orbital can accomadate a max of 2 electrons, 1 spin-up and 1 spin-down. If an orbital is full, we say that the electrons it holds are "spin-paired"
69
Aufbau Principle
Electrons occupy the lowest energy orbitals availbale (electron subshells are filled in order of increasing energy)
70
Hund's Rule
Electrons in the same subshell occupy available orbitals singly, before pairing up.
71
Pauli Exclusion Principle
There can be no more than 2 electrons in any given orbitals
72
Whats the max # of electrons that can go into any *s*, *p*, *d*, and *f* subshell?
s subshell: no more than 1 \* 2 = 2 electrons
p subshell: no more than 3 \*2 = 6 electrons
d subshell: no more than 5 \* 2 = 10 electrons
f subshell: no more than 7 \* 2 = 14 electrons
73
Noble gases in relation to their outer electrons
noble gases all have their 8 outer electron in filled subshells, 2 in the *ns* and 6 in the *np B*ecause there 8 valence electrons are in filled subshells, we say that these atoms have a complete octet which accounts for chemical inactivity.
74
Diamagnetic atoms
an atom that has **all of its electrons spin paired** is referred to as diamagnetic. A diamagnetic atom **must contain an even number** of electrons and have all of its occupied subshells filled. Atom will be **REPELLED** by externally produced magnetic field
75
Paramagnetic atoms
If atoms electrons are **not all spin paired**, it is said to be paramagnetic. Paramagentic atoms are **ATTRACTED** into externally produced magnetic fields.
76
Period
horizontal row in the periodic table
77
Group (family)
vertical column
78
Order of electron configurations
1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p
\*d block, subtract 1 from the period number
\*f block, subtract 2 from the period number
79
Anamolous Electron Configurations
atoms can achieve a lower energy state or higher degree of stability by having a filled or half filled d subshell.
\*Cr is rather 4s13d5 instead of 4s23d4
\*Cu is 4s13d10 rather than 4s23d9
\*Cr, Cu, Mo, Ag, Au are ALL exceptions that follow this rule
80
Atoms with the same electron configuration are said to be \_\_\_\_\_\_\_\_\_\_
isoelectronic
81
If an atom becomes an anion (-) (if it acquires more electrons) we move to the \_\_\_\_\_\_\_\_
RIGHT
82
If an atom becomes a cation (+) (if it loses more electrons) we move to the \_\_\_\_\_\_\_\_
LEFT
83
Electrons that are removed (ionized) from an atom always come from the \_\_\_\_\_\_\_\_
valence shell (the highest n level) and highest energy orbital within that level
Examples: Li to Li+ (1s22s1) becomes 1s2
Transtion Metals (elements in the d block have both ns and (n-1) d electrons. To form a cation (+) atoms will lose their valence electrons first and n is \> n-1, transiton metals lose s electrons before d electrons. Only after all s electrons are lost, do d get ionized
(Ti) has configuration of 4s23d2
Ti+ is [Ar] 4s13d2 and Ti2+ is [Ar] 3d2
84
Excited State vs. Ground State
just be sure to count # of electrons and make sure they stay the same and that one electron is jumped to another subshell
85
Valence Electrons
electrons in the outermost shell
86
Groups of the periodic table
* Group 1: Alkali Metals ns1
* Group 2: Alkali Earth Metals ns2
* Group 7: Halogens ns2np5
* Group 8: Noble Gases ns2np6
87
Special properties of the groups of the periodic table
filled valence shell config is called an octet and results in great stability. Most noble gases and helium is inert. Alkali metals and alkali earth metals behave as reducing agents (lose valence electrons). Halogens require only a single electron to achieve stable octet. Halogens naturally diatomic and behave as oxidizing agents (gain electrons)
88
Which elements are metalloids?
B, Si, Ge, As, Sb, Te, Po
89
nuclear shielding or the shielding effect
each filled shell between the nucleus and the valence electron shields or "protects" the valence electrons from the full effect of the positively charged proton in the nucleus. As far as the valence elctrons are concerned, the electrical pull by the protons in the nucleus is reduced by the negative charges of the electrons in the filled shells in between; the result is an effective reduction in the positive elementary charge from Z to a smaller amount denoted by Zeff (for effective nuclear charge)
90
Atomic Radius
As we go across a period, electrons are being added but new shells are not. Therefore the valence electrons are more and more tightly bound to the atom because they feel a greater effective nuclear charge. **So left to right across a period, atomic radius decreases.**
With progresion down a group, as new shells are added with each period, the valence electrons experience increased shielding. The valence electrons are less tightly bound since they feel a smaller effective nuclear charge. **As we go down a group, atomic radius increases due to increased shielding.**
91
Ionic Radius
If we form an ion, the radius will decrease as electrons are removed (because the ones that are left are drawn in more closesly to the nucleus) and the radius will increase as electrons are added. So in terms of radius, we have X+\< X -
92
Ionization Energy
The amount of energy neccesary to remove the least tightly bound electron from an isolated atom is called the atoms **first ionization energy .** As we move from left to right across a period and up a group, IE increases since valence electrons are more tightly bound. The IE of any atom with a noble gas config will always be very large. The **second ionization energy** of an atom X is the energy required to remove the least tightly bound lectron from the cation X+. **IE2 will always be greater than IE1**
93
Electron Affinity
The energy associated with the addition of an electron to an isolated atom is known as the atoms electron affinity. If energy is releases when the electron is added, we say EA is negative (-). If energy is required in order to add the electron, we say EA is positive (+) . Halogens have large negative EA values since the addition of an electron would give them an octet. Noble gases and alkali earth metals have positive EA values due to added electron fills a new lever or sublevel and destabilizes the electron config. GENERAL TREND is attached.
94
Electronegativity
F \> O \> N \> Cl \> Br \> I \> S \> C = H
95
Acidity
Acidity is measure of how well a compound donates protons, accepts electrons, or lowers PH in a chemical system
**Horizontal period trend for acidity:** the more electronegative the element is, the more stable the anion will be so acidity increases from left to right across a period
**Vertical trend for acidity:** depends on the size of the anion, the larger the anion, the more the negative charge will be localized and stabilized so acidity increases down a group in the periodic table.
