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Chemistry > Vocab Final Exam > Flashcards

Vocab Final Exam Flashcards

(107 cards)

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1
Q

Chemical bond

A

Mutual electrical attraction between nuclei and valence electrons of different atoms that binds the atoms together

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2
Q

Ionic bonding

A

Results from the electrical attraction between large numbers of cations and anions. atoms donate or accept electrons from each other

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3
Q

Covalent bonding

A

Results from sharing of electron pairs between 2 atoms. The electrons shared belong to both atoms

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4
Q

Polar covalent

A

When electrons are shared unevenly

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5
Q

Nonpolar covalent

A

When electrons are shared evenly

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6
Q

Ionic v. Covalent

A

Ionic: >1.7
Polar covalent: 0.3-1.7
Nonpolar covalent: 0-0.3
LOOK AT DIFFERENCE IN ELECTRONEGATIVITIES

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7
Q

Heat of sublimation

A

Amount of heat required to turn solid, bonded metal atoms into gaseous individual atoms

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8
Q

Properties of Metals

A

🔸conductivity
🔹luster(shininess)
🔸malleability & ductility

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9
Q

Octet Rule

A

A compound tends to form so that each atom has an octet(8) of electrons in its highest energy level by gaining, losing, or sharing electrons

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10
Q

Lewis dot structures

A

Like dot diagrams but for entire molecules, atomic symbols represent nucleus and core electrons and dots or dashes represent valence electrons

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11
Q

Double bonds

A

Sharing of 2 pairs of electrons; stronger and shorter

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12
Q

Triple bonds

A

Sharing of 3 pairs of electrons; strongest and shortest

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13
Q

Resonance

A

Bonding between atoms that cannnot be represented in on Lewis structure, shows all possible structures with double-ended arrow in between to show that electrons are delocalized

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14
Q

Single bond

A

Sharing of one pair of electrons; weakest and longest

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15
Q

Covalent network bonding

A

A different type of bonding, lots of nonmetal covalently bonded together in a network in all directions
Ex. Diamond, graphite, silicon dioxide

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16
Q

VSPER Theory

A

Valence, Shell, Electron, Pair, Repulsion
Repulsion between oaurs of electrons around an atom cause them to be as far apart as possible; used to predict the geometry of molecules

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16
Q

Intermolecular forces

A

Forces of attraction between molecules, vary in strength depending on type of molecule. Polar molecules have stronger IMF’s than nonpolar molecules. Are usually weaker than intramolecular forces( bonds inside a molecule)

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17
Q

Dipoles

A

A polar molecule( with an imbalance in electrons) has a dipole moment. There is a + center and - center
Represented by an arrow pointing from + center to - center

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18
Q

Boiling Point

A

Is a good measure of intermolecular forces(between molecules). Forces must overcome the attractive forces between molecules to pull them away from each other and enter the gas phase. Greater the molecules’s attraction, the higher the BP

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19
Q

Dipole-dipole

A

When molecules with dipole moments line up to minimize repulsion and maximize attraction. Very weak compared to covalent and ionic bonds

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20
Q

Hydrogen Bonding

A

🔸happens between H, N, O, & F

🔹 very strong type of dipole-dipole attraction because bond is so polar and because atoms are so small

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21
Q

London dispersion forces

A

In every molecular compound, only important for nonpolar molecules and noble gas atoms
🔹 weak, short-lived
🔸caused by formation of temporary dipole moments

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22
Q

Ionic crystals

A

🔹sites in lattice contain ions
🔸hard, brittle
🔹high melting points
🔸 good insulators as solids

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23
Q

Subscripts

A

Give you the simplest ratio of atoms or the number of atoms in a molecule

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24
Molar mass
Mass of one mole of pure substance, units: g/mol
25
Percent composition
Percentage of mass of each element in the compound | Mass of the element in sample/ total mass of sample • 100
26
Chemical equation
Represents identities and relative amounts of reactants and products in the chemical reaction
27
Evidence or chemical change
🔸heat or light 🔹gas bubbles 🔸precipitate 🔹color change
28
Word equations
Uses names instead of formulas
29
Coefficients
Specifies the relative number of moles and molecules involved in the reaction, used to balance the equation
30
Synthesis
Combining more than one reactant to make only one single product
31
Decomposition
Breaking apart one reactant to make more than one product
32
Single replacement
An element replaces a similar element in a compound
33
Double Replacement
Two similar elements switch places
34
Combustion
Combines with oxygen and releases energy in form of heat/ light 🔹H2O and usually CO2 as products
35
Limiting reactant
Reactant that limits the amount of product created; completely used up
36
Excess reactant
React not completely used up
37
Theoretical yield
Maximum amount of product possible
38
Actual yield
The measured amount formed in reaction
39
Percent yield equation
Actual yield/ theoretical yield • 100
40
Activity series
Used to predict whether single replacement reactions will occur, most actce is on top, an element can replace anything below it but not any above it
41
Kinetic-molecular theory
Based on the idea that particles of matter are always in motion, theory can be used to explain the properties of solids, liquids, & gases in terms of energy of particles and the forces that act between them
42
Ideal gas
Is a hypothetical gas that perfectly fits all the assumptions of KMT
43
Elastic collision
Is one in where there is no net loss of total kinetic energy
44
Fluids
Gases and liquids flow, therefore they are fluids
45
Real gas
Is a gas that doesn't behave completely according to the assumptions of KMT
46
Surface tension
Force that pulls parts of a liquid's surface together
47
Capillary action
Attraction of surface of liquid to surface of a solid Ex. Water & paper Water & glass Meniscus
48
Enthalphy of vaporization
Energy required to vaporize 1 mole of a liquid at 1 atm; ENDOTHERMIC
49
Freezing
Physical change of a liquid to solid by removal of heat
50
Enthalphy of fusion
Heat required to be removed between liquid to a solid
51
Vapor pressure
Liquid can be injected under inverted tube, part of the liquid evaporates to the top of Pvap can be determind by height of Hg(mercury)
52
Volatile
Liquids with high vapor pressure, evaporate quickly, weak intermolecular forces
53
Phase diagrams
Represents the phases as function of T and P
54
Triple point
All 3 phases exist- all have same vapor pressure
55
Critical point
Where if the temperature is increased, vapor can't be liquefied no matter what pressure is applied. Only one fluid phase exsists that is called supercritical fluid
56
Pressure equation
Force/ area
57
Barometer
A device used to measure atmospheric pressure
58
Dalton's law
The total pressure of a gas mixture is the sum of the partial pressures of the component gases Ptotal= Pvap + Pgas COLLECTING GAS OVER WATER
59
Boyle's law
The V of fixed mass of gas varies inverseky with P at a constant T P1V1= P2V2
60
Charles's Law
The V of fixed mass of gas at constant P varies directly with Kelvin T V1/T1 = V2/T2
61
Gay-Lussac's Law
The P of fixed mass of gas at constant V varies directly with Kelvin T P1/T1 = P2/T2
62
Combined gas law
Relationship between P, V, and T of a fixed amount of gas | P1V1/T1 = P2V2/T2
63
Avogadro's law
The law states that equal volumes of gases at the same temperature and pressure contsin equal numbers of molecules V=kn or V/n =k
64
Standard molar volume
The volume occupied by one mole of gas at STP or 22.4L
65
Ideal gas law
Mathematical relationship between pressure, volume, temperature, and number of moles of a gas PV=nRT
66
Ideal gas constant
Constant "R" = 0.0821 L•atm/ mol•K= 8.314 L•kPA/ mol•K= 62.4 L•mmHg/ mol•K
67
Effusion
When gases escape through a tiny opening into a vaccum
68
Diffusion
When gases mix automatically without being stirred
69
Graham's Law of Effusion
Rates of effusion of gases at a constant T and P are inverseky proportional to the square roots of their molar masses Rate of effusionA/ Rate of effusionB = square root MB/MA
70
Mole ratio
of moles of one substance to ratio of # of moles of another substance
71
Sublimation
Solid to gas
72
Condensation
Gas to liquid
73
Vaporization
Liquid to gas
74
Deposition
Gas to solid
75
Boiling
Example of vaporization
76
Freezing
Liquid to solid
77
Melting
Solid to liquid
78
Cations
Positive ions
79
Anions
Negative ions
80
Melting point
Solid and liquid phase exsist in equilibrum
81
Freezing point
Water; 0C
82
Standard Temperature and Pressure(STP)
Pressure: 1atm Volume: 22.4L Temperature: 273.15K
83
Formula unit
Empirical formula of any ionic or covalent network solid compound used as an independent entity for stoichionetric calculations. Lowest whole number ratio of ions represented in an ionic compound
84
Polyatomic ion
Is a charged chemical ion composed of 2 or more atoms covalently bonded
85
Hybridization
Mixing atomic orbitals ino bee hybrid orbitals with different energies, shapes, etc. than the component atomic orbitals
86
Binary compound
Is a chemical compound that contains exactly 2 different types of elements Ex. Calcium chloride, sodium fluoride
87
Oxyanion
Is an ion with the generic formua A xOz-y. Determined by the octet rule
88
Monatomic ion
Is an ion consisting of a single atom
89
Salt
NaCl; remaining solid is known as the anhydrous salt
90
Oxidation number
Indicator of the degree of oxidation( loss of electrons) of an atom in a chemical compound
91
Oxidation state
Hypothetical charge that an atom would have if all bonds to atoms of different elementd were 100% ionic, with no covalent component
92
Precipitate
Solid formed when the reaction occurs in a liquid solution
93
Reversible reaction
Chemical reaction where the reactants form products that, in turn, react together to give the reactants back
94
Percent yield
Amount of product obtained in a chemical reaction | Actual yield/ theoretical yield •100
95
Dissociation
General process in which molecules separate or split into smaller particles such as atoms, ions, or radicals
96
Ionization
Any process by which electrically neutral atoms or molecules are converted to ions
97
Net ionic equation
A chemical equation for a reaction which lists only those species participating in the reaction
98
Spectator ion
Ion that exsists as a reactsnt and product in a chemical reaction
99
Suspension
Heterogeneous mixture containing solid particles that are sufficiently large for sedimentation
100
Molarity
Number of moles of solute per liter of solution
101
Molality
Number of moles of soluter per kilogram of solvent
102
Mole fraction
To be equal to the number of moles of a component divided by the total number of moles of a solution
103
Mass percent
Shows the amount each element in a molecule contributed to the total molecular mass
104
Atmospheric pressure
Pressure exerted by the weight of air in the atmosphere of Earth
105
Manometer
Instrument for measuring the pressure acting on a column of fluid
106
Amorphous solid
A solid which doesn't exhibit crystalline structure. | Ex. Window glass, polystyrene

Chemistry (8 decks)

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