Gre Chemistry Flashcards
1
Q
Mass number
A
Total number if protons and neutrons in the nucleus represented by letter A
2
Q
Atomic number
A
Total number if protons in the nucleus represented by letter Z
3
Q
Isotope
A
Two or more nuclei of the same element that have different mass numbers
4
Q
Binding energy
A
Energy required to overcome proton proton repulsion and hold the nucleus together
5
Q
Parent nucleus
A
Nucleus prior to nuclear decay
6
Q
Daughter nucleus
A
Nucleus formed as a result of nuclear decay
7
Q
Alpha decay
A
Helium nuclei 2 protons 2 neutrons
8
Q
Beta decay
A
Releases electron daughter will always be different element but will have the same mass
9
Q
Gamma decay
A
Daughter is identical to parent except it has less energy
10
Q
First order decay
A
Probability that a nucleus will decay in a given time is constant and independent of surroundings.
the rate of loss of mass at any given time is directly proportional to the mass present at that time.
11
Q
Nuclear fission
A
When an accelerated particle such as a neutron striking a nuclei the nucleus can split into two or more fragments
12
Q
Nuclear fusion
A
Accelerated particle captured by nucleus to create larger nucleus
13
Q
Bohr atom
A
Atom with only one electron
14
Q
Energy shells
A
Certain distances away from the nucleus
15
Q
Ground state
A
Configuration where electrons are in lowest energy levels
16
Q
Excited state
A
Configuration where electrons are not in lowest energy levels
17
Q
Valences electrons
A
Electrons with the largest value of n
18
Q
Principle quantum number
A
n, energy level > or equal to n
19
Q
Secondary quantum number
A
I, angular momentum, 0-n
20
Q
Degenerate
A
Quantum states or configurations that have identical energies
21
Q
Magnetic quantum number
A
quantum number that identifies different orbitals within a subshell. ml can take on values from -l to +l. The number of orbitals within a subshell is the number of possible magnetic quantum number values.
22
Q
S
A
1 orbital
23
Q
P
A
3 orbitals
24
Q
D
A
5 orbitals
25
F
7 orbitals
26
Spin quantum number
The fourth quantum number denoted by ms. The spin quantum number indicates the orientation of the intrinsic angular momentum of an electron in an atom. The only possible values of a spin quantum number are +½ or -½ (sometimes referred to as 'spin up' and 'spin down').
27
Pauli exclusion principle
The Pauli exclusion principle states no two electrons can have the identical quantum mechanical state in the same atom.
No pair of electrons in an atom can have the same quantum numbers n, l, ml and ms.
28
Aufbau principle
no two electrons in the atom will share the same four quantum numbers n, l, m, and s.
electrons will first occupy orbitals of the lowest energy level.
electrons will fill an orbital with the same spin number until the orbital is filled before it will begin to fill of the opposite spin number.
electrons will fill orbitals by the sum of the quantum numbers n and l. Orbitals with equal values of (n+l) will fill with the lower n values first.
29
Hunds rule
When partially filling degenerate orbitals of p and d and f always put one electron in each orbital before pairing them up. Orient them so their magnetic spins are all the same
30
Atomic size trend
Increases as one moves down or to the left in the periodic table
31
Atomic and ionic radii
Cation < neutral < anion
32
Isoelectronic species
Two atoms or ions that have the same number of electrons
33
Ionization
Removal of an electron from a neutral atom
34
Endergonic process
Requires input of energy
35
Ionization energies
The first ionization energy is the energy required to remove one electron from the parent atom. The second ionization energy is the energy required to remove a second valence electron from the univalent ion to form the divalent ion, and so on. Successive ionization energies increase. The second ionization energy is always greater than the first ionization energy.
36
Electron affinity
Electron affinity reflects the ability of an atom to accept an electron. It is the energy change that occurs when an electron is added to a gaseous atom. Atoms with stronger effective nuclear charge have greater electron affinity.
37
Exergonic
Exergonic refers to a chemical reaction where the free energy of the system decreases.
38
Electro negativity
Electronegativity is related to ionization energy. Electrons with low ionization energies have low electronegativities because their nuclei do not exert a strong attractive force on electrons. Elements with high ionization energies have high electronegativities due to the strong pull exerted on electrons by the nucleus.
Increases as you move up and to the right
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Acidity
Increases as one moves down and to the right
40
Basicity
Increases as one moves up and to the left
41
Octet
An eight electron arrangement in the outer electron shell of the noble gases
42
Cation
Positively charged ion
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Anion
Negatively charged ion
44
Metalloids
Metalloids are located along the line between the metals and nonmetals in the periodic table. The metalloids are boron, silicon, germanium, arsenic, antimony, and tellurium. Polonium
45
Properties of metalloids
The reactivity of the metalloids depends on the element with which they are reacting. For example, boron acts as a nonmetal when reacting with sodium yet as a metal when reacting with fluorine. The boiling points, melting points, and densities of the metalloids vary widely. The intermediate conductivity of metalloids means they tend to make good semiconductors.
46
Valence shell electron pair repulsion
Electron pairs wether bonding or non bonding tend to move as far apart as possible
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Non polar bond
Type of chemical bond which has no positive or negative 'ends'
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Malleability
capable of being shaped. The term is often used with reference to metals, as in the degree to which they can be shaped by pounding with a hammer.
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Ductility
Ductile is a physical property of a material associated with their ability to be hammered thin or stretched into wire without breaking.
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Band theory
Metal is thought of as a giant molecule in which delocalized molecular orbitals cover the entire structure
51
Intrinsic semi conductor
Solid in which the band gap is so small that some electrons from the valence band will occupy energy levels in the conduction band
52
P type conductivity
Charge carriers are positive holes
53
Coordinate covalent bonds
A coordinate bond is a covalent bond between two atoms where one of the atoms provides both electrons that form the bond
54
Coordination complex
A coordination complex or metal complex is a chemical species consisting of a central atom or ion bonded to surrounding molecules or ions. The central atom of a coordination complex commonly is a metal cation. Various ligands or complexing agents may surround the central atom of a coordination complex.
55
Covalent bond
A covalent bond is a chemical link between two atoms in which electrons are shared between them
56
Lewis base
A Lewis base is an substance that is an electron pair donator
57
Lewis acid
Molecule that accepts a pair of electrons to form a covalent bond
58
Intermolecular forces
The intermolecular force is the sum of all the forces between two neighboring molecules.
59
Intramolecular forces
The intramolecular force is the sum of all the forces holding a molecule or compound together.
60
Electron cloud repulsion
Prevents atoms from passing right through one another
61
Ionic bond
An ionic bond is a chemical link between two atoms caused by the electrostatic force between oppositely-charged ions in an ionic compound.
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Polar bond
A polar bond is a covalent bond between two atoms where the electrons forming the bond are unequally distributed. This causes the molecule to have a slight electrical dipole moment where one end is slightly positive and the other is slightly negative.
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Polar molecule
Have an asymmetric distribution of electron density
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Dipole forces
Interactions between the charges portions of polar molecules
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London dispersion force
Weak intermolecular force between two atoms or molecules in close proximity of each other. The force is a quantum force generated by electron repulsion between the electron clouds of two atoms or molecules as they approach each other.
66
Covalent compound
a molecule formed by covalent bonds, in which the atoms share one or more pairs of valence electrons
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Ionic compounds
properties of ionic compounds relate to how strongly the positive and negative ions attract each other in an ionic bond.
68
Unit cell
A unit cell is the smallest unit of a crystal, which, if repeated, could generate the whole crystal.
69
Ionic solids
+ and - ions arranged in regular arrays each ion is surrounded by ions of the opposite charge.
Hard brittle have high melting / boiling points
Poor conductors
70
Covalent solids
Made of atoms held together by strong covalent bonds. Very hard have high melting points
71
Molecular solids
A molecular solid is a type of solid in which molecules are held together by van der Waals forces rather than by ionic or covalent bonds.
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A buffer
A buffer is an aqueous solution that has a highly stable pH. If you add acid or base to a buffered solution, its pH will not change significantly. Similarly, adding water to a buffer or allowing water to evaporate will not change the pH of a buffer.
73
Chelate
A chelate is an organic compound formed when a polydentate ligand bonds to a central metal atom.
74
Hydroxyl group
functional group consisting of a hydrogen atom covalently bonded to an oxygen atom.
denoted by -OH in chemical structures and has a valence charge of -1.
75
Carboxyl group
organic functional group consisting of a carbon atom double bonded to an oxygen atom and single bonded to a hydroxyl group
commonly written as -C(=O)OH or -COOH
76
Alkali metals
lithium, sodium, potassium, rubidium, cesium, and francium.
exhibit many of the physical properties common to metals, although their densities are lower than those of other metals. have one electron in their outer shell, which is loosely bound. This gives them the largest atomic radii of the elements in their respective periods
77
Le chateliers principle
If the temperature, concentration, volume, or partial pressure of a chemical system at equilibrium changes, then the equilibrium shifts to compensate for the change.
78
Lattice energy
ΔH (enthalpy change) for the process in which oppositely charged ions in the gas phase combine to form an ionic lattice in the solid phase.
79
Hydrocarbon
hydrocarbon is a substance consisting only of carbon and hydrogen atoms.
Benzene hexane
80
Alkane
An alkane is a hydrocarbon containing only single carbon-carbon bonds.
The general formula for an alkane is CnH2n+2 where n is the number of carbon atoms in the molecule.
81
Non metals
hydrogen, carbon, nitrogen, oxygen, phosphorus, sulfur, and selenium.
Nonmetals have high ionization energies and electronegativities. They are generally poor conductors of heat and electricity. Solid nonmetals are generally brittle, with little or no metallic luster. Most nonmetals have the ability to gain electrons easily
82
Halogen
fluorine, chlorine, bromine, iodine, astatine, and ununseptium
have seven valence electrons. As a group, halogens exhibit highly variable physical properties. Halogens range from solid (I2) to liquid (Br2) to gaseous (F2 and Cl2) at room temperature
High electronegativities
83
Metals
Most of the elements on the periodic table are metals, including gold, silver, platinum, mercury, uranium, aluminum, sodium and calcium. Alloys, such as brass and bronze, also are metals.
Metals are shiny solids are room temperature (except mercury, which is a shiny liquid element), with characteristic high melting points and densities. Many of the properties of metals, including large atomic radius, low ionization energy, and low electronegativity, are due to the fact that the electrons in the valence shell of a metal atoms can be removed easily.
84
Alkali metals
lithium, sodium, potassium, rubidium, cesium, and francium
Lower densities than other metals
One loosely bound valence electron
Largest atomic radii in their periods
Low ionization energies
Low electronegativities
Highly reactive
85
Ionization energy
The ionization energy, or ionization potential, is the energy required to completely remove an electron from a gaseous atom or ion. The closer and more tightly bound an electron is to the nucleus, the more difficult it will be to remove, and the higher its ionization energy will be.
Ionization energy is measured in electronvolts (eV). Sometimes the molar ionization energy is expressed, in J/mol.
86
Alkaline earths
They have smaller atomic radii than the alkali metals. The two valence electrons are not tightly bound to the nucleus, so the alkaline earths readily lose the electrons to form divalent cations.
87
Transition metals
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Low ionization energies
Positive oxidation states
Very hard
High melting points
High boiling points
High electrical conductivity
Malleable
Five d orbitals become more filled, from left to right on periodic table
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88
Solubility
he maximum quantity of a substance that may be dissolved in another. The maximum amount of solute that may be dissolved in a solvent.
89
Solute
The substance that is dissolved in a solution. For solutions of fluids, the solvent is present in greater amount than the solute.
90
Solvent
The component of a solution that is present in the greatest amount. It is the substance in which the solute is dissolved.
: The solvent for seawater is water. The solvent for air is nitrogen
91
Mixture
Two or more substances which have been combined such that each substance retains its own chemical identity.
92
First law of thermodynamics
he law which states that the total energy of a system and its surroundings remains constant.
93
Third law of thermodynamics
The third law of thermodynamics states the entropy of a perfect crystal approaches zero as the temperature approaches absolute zero.
94
Thermal contact
When two substances can effect each others temperature
95
Thermal Equilibrium
is when two substances in thermal contact no longer transfer heat.
96
Thermal expansion
when a substance expands in volume as it gains heat. Thermal contraction also exists.
97
Conduction
When heat flows through a heated solid.
98
Convection
when heated particles transfer heat to another substance, such as cooking something in boiling water.
99
Radiation
when heat is transferred through electromagnetic waves, such as from the sun
100
Insulation
when a low-conducting material is used to prevent heat transfer.
101
Thermodynamic processes
Adiabatic process - a process with no heat transfer into or out of the system.
Isochoric process - a process with no change in volume, in which case the system does no work.
Isobaric process - a process with no change in pressure.
Isothermal process - a process with no change in temperature.
102
Condensation
Gas to liquid
103
Freezing
Liquid to solid
104
Melting
Solid to liquid
105
Sublimation
Solid to gas
106
Vaporization
Liquid or solid to gas
107
Heat capacity
C, of an object is the ratio of change in heat (energy change - denoted by delta-Q) to change in temperature (delta-T).
C = delta-Q / delta-T
indicates ease with which a substance heats up.
108
Second law of thermodynamics
It is impossible for a process to have as its sole result the transfer of heat from a cooler body to a hotter one.
109
Zeroth law of thermodynamics
Two systems each in thermal equilibrium with a third system are in thermal equilibrium to each other.
110
Thermal radiation
electromagnetic radiation emitted by objects because of their temperature
111
Ph
measure of hydrogen ion concentration; a measure of the acidity or alkalinity of a solution. Aqueous solutions at 25°C with a pH less than seven are acidic, while those with a pH greater than seven are basic or alkaline. A pH level of is 7.0 at 25°C is defined as 'neutral
112
Oxidation
Loss of hydrogen loss of electrons gain hydrogen
113
Reduction
Gain of hydrogen gain electrons loss oxygen
