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Gre Chemistry Flashcards

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1
Q

Mass number

A

Total number if protons and neutrons in the nucleus represented by letter A

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2
Q

Atomic number

A

Total number if protons in the nucleus represented by letter Z

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3
Q

Isotope

A

Two or more nuclei of the same element that have different mass numbers

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4
Q

Binding energy

A

Energy required to overcome proton proton repulsion and hold the nucleus together

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5
Q

Parent nucleus

A

Nucleus prior to nuclear decay

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6
Q

Daughter nucleus

A

Nucleus formed as a result of nuclear decay

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7
Q

Alpha decay

A

Helium nuclei 2 protons 2 neutrons

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8
Q

Beta decay

A

Releases electron daughter will always be different element but will have the same mass

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9
Q

Gamma decay

A

Daughter is identical to parent except it has less energy

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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.

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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

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12
Q

Nuclear fusion

A

Accelerated particle captured by nucleus to create larger nucleus

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13
Q

Bohr atom

A

Atom with only one electron

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14
Q

Energy shells

A

Certain distances away from the nucleus

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15
Q

Ground state

A

Configuration where electrons are in lowest energy levels

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16
Q

Excited state

A

Configuration where electrons are not in lowest energy levels

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17
Q

Valences electrons

A

Electrons with the largest value of n

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18
Q

Principle quantum number

A

n, energy level > or equal to n

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19
Q

Secondary quantum number

A

I, angular momentum, 0-n

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20
Q

Degenerate

A

Quantum states or configurations that have identical energies

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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.

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22
Q

S

A

1 orbital

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23
Q

P

A

3 orbitals

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24
Q

D

A

5 orbitals

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25
Q

F

A

7 orbitals

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26
Q

Spin quantum number

A

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’).

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27
Q

Pauli exclusion principle

A

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.

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28
Q

Aufbau principle

A

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.

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29
Q

Hunds rule

A

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

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30
Q

Atomic size trend

A

Increases as one moves down or to the left in the periodic table

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31
Q

Atomic and ionic radii

A

Cation < neutral < anion

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32
Q

Isoelectronic species

A

Two atoms or ions that have the same number of electrons

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33
Q

Ionization

A

Removal of an electron from a neutral atom

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34
Q

Endergonic process

A

Requires input of energy

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35
Q

Ionization energies

A

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.

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36
Q

Electron affinity

A

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.

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37
Q

Exergonic

A

Exergonic refers to a chemical reaction where the free energy of the system decreases.

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38
Q

Electro negativity

A

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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39
Q

Acidity

A

Increases as one moves down and to the right

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40
Q

Basicity

A

Increases as one moves up and to the left

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41
Q

Octet

A

An eight electron arrangement in the outer electron shell of the noble gases

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42
Q

Cation

A

Positively charged ion

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43
Q

Anion

A

Negatively charged ion

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44
Q

Metalloids

A

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

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45
Q

Properties of metalloids

A

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.

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46
Q

Valence shell electron pair repulsion

A

Electron pairs wether bonding or non bonding tend to move as far apart as possible

47
Q

Non polar bond

A

Type of chemical bond which has no positive or negative ‘ends’

48
Q

Malleability

A

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.

49
Q

Ductility

A

Ductile is a physical property of a material associated with their ability to be hammered thin or stretched into wire without breaking.

50
Q

Band theory

A

Metal is thought of as a giant molecule in which delocalized molecular orbitals cover the entire structure

51
Q

Intrinsic semi conductor

A

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
Q

P type conductivity

A

Charge carriers are positive holes

53
Q

Coordinate covalent bonds

A

A coordinate bond is a covalent bond between two atoms where one of the atoms provides both electrons that form the bond

54
Q

Coordination complex

A

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
Q

Covalent bond

A

A covalent bond is a chemical link between two atoms in which electrons are shared between them

56
Q

Lewis base

A

A Lewis base is an substance that is an electron pair donator

57
Q

Lewis acid

A

Molecule that accepts a pair of electrons to form a covalent bond

58
Q

Intermolecular forces

A

The intermolecular force is the sum of all the forces between two neighboring molecules.

59
Q

Intramolecular forces

A

The intramolecular force is the sum of all the forces holding a molecule or compound together.

60
Q

Electron cloud repulsion

A

Prevents atoms from passing right through one another

61
Q

Ionic bond

A

An ionic bond is a chemical link between two atoms caused by the electrostatic force between oppositely-charged ions in an ionic compound.

62
Q

Polar bond

A

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.

63
Q

Polar molecule

A

Have an asymmetric distribution of electron density

64
Q

Dipole forces

A

Interactions between the charges portions of polar molecules

65
Q

London dispersion force

A

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
Q

Covalent compound

A

a molecule formed by covalent bonds, in which the atoms share one or more pairs of valence electrons

67
Q

Ionic compounds

A

properties of ionic compounds relate to how strongly the positive and negative ions attract each other in an ionic bond.

68
Q

Unit cell

A

A unit cell is the smallest unit of a crystal, which, if repeated, could generate the whole crystal.

69
Q

Ionic solids

A

+ 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
Q

Covalent solids

A

Made of atoms held together by strong covalent bonds. Very hard have high melting points

71
Q

Molecular solids

A

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.

72
Q

A buffer

A

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
Q

Chelate

A

A chelate is an organic compound formed when a polydentate ligand bonds to a central metal atom.

74
Q

Hydroxyl group

A

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
Q

Carboxyl group

A

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
Q

Alkali metals

A

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
Q

Le chateliers principle

A

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
Q

Lattice energy

A

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
Q

Hydrocarbon

A

hydrocarbon is a substance consisting only of carbon and hydrogen atoms.

Benzene hexane

80
Q

Alkane

A

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
Q

Non metals

A

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
Q

Halogen

A

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
Q

Metals

A

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
Q

Alkali metals

A

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
Q

Ionization energy

A

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
Q

Alkaline earths

A

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
Q

Transition metals

A 
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
88
Q

Solubility

A

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
Q

Solute

A

The substance that is dissolved in a solution. For solutions of fluids, the solvent is present in greater amount than the solute.

90
Q

Solvent

A

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
Q

Mixture

A

Two or more substances which have been combined such that each substance retains its own chemical identity.

92
Q

First law of thermodynamics

A

he law which states that the total energy of a system and its surroundings remains constant.

93
Q

Third law of thermodynamics

A

The third law of thermodynamics states the entropy of a perfect crystal approaches zero as the temperature approaches absolute zero.

94
Q

Thermal contact

A

When two substances can effect each others temperature

95
Q

Thermal Equilibrium

A

is when two substances in thermal contact no longer transfer heat.

96
Q

Thermal expansion

A

when a substance expands in volume as it gains heat. Thermal contraction also exists.

97
Q

Conduction

A

When heat flows through a heated solid.

98
Q

Convection

A

when heated particles transfer heat to another substance, such as cooking something in boiling water.

99
Q

Radiation

A

when heat is transferred through electromagnetic waves, such as from the sun

100
Q

Insulation

A

when a low-conducting material is used to prevent heat transfer.

101
Q

Thermodynamic processes

A

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
Q

Condensation

A

Gas to liquid

103
Q

Freezing

A

Liquid to solid

104
Q

Melting

A

Solid to liquid

105
Q

Sublimation

A

Solid to gas

106
Q

Vaporization

A

Liquid or solid to gas

107
Q

Heat capacity

A

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
Q

Second law of thermodynamics

A

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
Q

Zeroth law of thermodynamics

A

Two systems each in thermal equilibrium with a third system are in thermal equilibrium to each other.

110
Q

Thermal radiation

A

electromagnetic radiation emitted by objects because of their temperature

111
Q

Ph

A

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
Q

Oxidation

A

Loss of hydrogen loss of electrons gain hydrogen

113
Q

Reduction

A

Gain of hydrogen gain electrons loss oxygen

Chemistry (2 decks)

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