Chem Ch. 5.3-6 Flashcards Preview

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Flashcards in Chem Ch. 5.3-6 Deck (40):
1

Mole

Is an amount of any substance or item that contains the same number of elementary units as there are atoms in exactly 12g of carbon-12.

2

Formula mass

The sum of masses of the atoms represented in the formula

3

Molar mass

The mass of one mole of that substance.

The molar mass is numerically equal to the atomic mass, molecular mass, or formula mass, but it is expressed in the unit grams per mole. (g/mol).

Ex: atomic mass of sodium 23.0 u, son it’s molar mass is 23.0 g/mol; the molecular mass of carbon dioxide is 44.0 u, its molar mass is 44.0 g/mol; the formula mass for ammonium sulfate is 132.1 u, so its molar mass is 132.1 g/mol

1 mol Na = 23.0 g Na
1 mol CO2 = 44.0 g CO2
1 mol (Nh4)2SO4 = 132.1 g (Nh4)2SO4

4

Avogadro’s number

6.02 x 10^23

5

Stoichiometry

The quantitative relationship between reactants and products in a chemical reaction.

6

Solution

A homogeneous mixture of two or more substances.

The substance being dissolved is the solute, and the substance doing the dissolving is the solvent.

7

Molarity

The amount of solute, in moles, per liter of solution.

Molarity (M)= moles of solute / liters of solution

8

Solubility

Maximum quantity of a given solute that can dissolve in any given kind of solvent at a specific temperature. (Solubility can increase, stay the same, or decrease with an increase in temperature depending upon the substance in question.)

9

Insoluble

A substance that can not be dissolved in any given solvent (potentially imprecise term.)

10

mM

The amount of dissolved substance in units of one thousandth of a mole per liter.

11

MW

Molecular weight/molar mass

Mass in grams of one mole of any substance

12

Valence electrons

Electrons in the outermost shell.

13

Electronegativity
EN

The relative tendency of an atom to attract or draw electrons to itself.

14

Ionic compounds

Compounds composed of positive and negative ions and held together by ionic bonds.

Ionic bonds- Strong chemical bond formed by the attraction between two oppositely charged ions. ie: between positively charged cation and negatively charged anion.
Cation- lost 1 or more electrons
Anion- gained 1 or more electrons

15

Formula unit

Smallest repeating unit of a salt.

16

Salt

Ions must be paired with oppositely charged ions. The resulting electrically neutral compound is called a salt.

Salts form by complete transfer of electrons from a cation to anion

17

Polyatomic ion

A positive or negative ion composed of two or more atoms.

*Polyatomic ion formula enclosed in parentheses when more than one is present

18

Non-bonding pairs

Lone pairs

Valence electrons in a molecule that aren’t in a bond between atoms; have significance in chemical reactions, so need to be noted.

19

“-ide”

Names of simple negative ions (anions) are derived from those of their parent elements by replacing the usual ending with a suffix “-ide” And adding the word ion.

Ex: when chlorine atom gains an electron, it becomes a chloride ion (Cl^-).

20

Free radical

And atom or molecule that contains an unpaired electron

21

Hydrocarbons

A class of organic molecules that contain only carbon and hydrogen.

(methane is simplest hydrocarbon possible).

22

Alcohols

A class of organic molecules related to hydrocarbons where at least one hydrogen was replaced by a hydroxide ion.

(Small alcohols tend to be more polar than non-polar, so are soluble in water.)

23

Metallic bonding

Occurs between metals.
- metals achieve a noble gas electron configuration by losing electrons, but without an electron receptor.
- Electrons formed by ionization of a metal “float around” in a “sea of charge” and neutralize positive charge of individual metal ions.

24

“-per”

Prefix meaning one more oxygen than the -ate.
Hypo-
-ite
-ate
Per-

25

“Hypo-“

Prefix meaning one less oxygen than the -ite.
Hypo-
-ite
-ate
Per-

26

Vapor pressure

The increase in pressure above a liquid caused by molecules of that substance that have escaped from the liquid. (relative humidity is derived from vapor pressure)

Vapor pressure increases with increasing temperature.

27

Boiling point

Temperature at which vapor pressure of a liquid equals atmospheric pressure.

28

Volatile

Easily vaporized substances.

29

Evaporation (related to vaporization)

Individual molecules at surface of a liquid acquire enough energy to escape; slower than boiling; can occur at temperatures well below boiling point.

30

Condensation

The change in the physical state of a substance from the gaseous phase into a liquid.

31

Melting point

Temperature at which a solid turns into a liquid.

32

Dipole

Polar covalent molecules have partial positive and negative ends that are attracted to one another. Example: water

33

Hydrogen bonding

When a hydrogen atom covalently bonded to a strongly electronegative atom with non-bonding electron pairs. Main elements where this occurs: O, N and F. H will lose a little of its negative charge to bond with O, N & F.

34

Solutions

Homogenous mixtures of two or more substances, often a liquid one.

35

Hydrated

When water molecules pack around each ion as tightly as possible. => ions become hydrated. Process known as hydration.

36

Specific heat capacity

The amount of heat required to raise 1g of a substance 1 degree celsius

37

Saturated

When limit reached, the water has absorbed all the ions it can and the solution has become saturated.

At the point of saturation, salt crystals will appear at the bottom of the solution.

38

Precipitation

Condensation of a solid from a solution during a chemical reaction.

Dissolved substance becoming a solid.

39

Miscibility

Ability to mix with other gases in any proportions; i.e., there is always room for more molecules of any kind.

Gases mix due to random molecular motion.

40

Sublimation

The transition from the solid phase to the gas phase without passing through an intermediate liquid phase.

This endothermic phase transition occurs at temperatures and pressures below the triple point.