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Flashcards in CH. 8 Deck (55):
1

Define equivalence point

equivalence of acid solution is = to equivalence of base.

2

Define end point

When indicator changes color

3

Define indicator

pH meter; used to determine if solution is at equivalence point.

4

Define molarity (M)

Moles of solute per L. Of solution

5

Define Neutralization

Acid and base react equivalently w/each other

6

Define titration

Solution of known concentration used to determine concentration of unknown solution

7

Define titrant

Known solution in titration

8

Define analyze

Unknown solution in titration

9

Define solute

Component in solution other than solvent (least amount in solution)

10

Define solution

Homogeneous mixture. Often liquid but not always

11

Define solvent

Largest component of solution (most moles present in solution)

12

Define solubility

Maximum quantity of substance that can dissolve in given volume of solution

13

Water soluble rules (which ones are they) and charges

Nitrates, A1 group, acetates, and ammonium.
Nitrates: -
A1: +
Acetates: -
Ammonium: +

14

Concentration units of % and M equation

M1V1=M2V2

C1V1=C2V2
(Dilution)

15

Precipitation equation

1. mL of 1st compound • (M1/1000mL)•(1 mol broken down compound/1 mol 2nd broken down compound)
2. mL of 2nd compound • (M2/1000mL)•(1 mol broken down compound/1 mol 2nd broken down compound)
3. Limiting agent •(molar mass of product/1 mol of product)

*on limiting agent attach elemental compound name

16

Define electrolyte

A solute that produces ions in solution which enables the solution to conduct electricity

17

Define strong electrolyte

Nearly 100% of dissociated into ions, conducts current efficiently, any solutions of ionic compounds strong acids and bases
Ex: NaCl, HCl

18

Define weak electrolyte

Slightly conductive, only partially dissociate into ions
Ex: tap water, weak acids and bases

19

Define nonelectrolyte

Substances in which no ionization occurs; no conduction of electrical current
Ex: solid NaCl, ethanol

20

Arrhenius acid/base define

Acid:
Substance when dissolved in water, produces hydronium (H3O+)
Base:
substance that, when dissolved in water produces hydroxide (OH-)

21

Bronsted-Lowery acid/base define

Acid:
Is a proton donor, must have removable proton
Base:
Is a proton acceptor, must have pair of nonbonding electrons

22

Strong acids

HCl, HBr, HI, H2SO4, HNO3, HClO4

23

Strong bases

1A group, 2A group, anything with OH

24

Titration calculation

V1M1/n1=V2M2/n2

n1=# of moles acid
n2=# of moles base

25

Properties of gases

•neither definite shape nor definite volume
•Uniformly fills any container
•exert pressure on surroundings
•volume changes with temperature and pressure •Mixes completely with other gasses and less dense than solids and liquids

26

Kinetic Theory of gas

•ransoms movement at high velocities
• no attractive/repulsive forces with other molecules
• in constant motion, moving rapidly in straight lines
•occupies larger volume than volume of other molecules
• engages in elastic collisions with wall of container and other gas molecules
•Has average kinetic energy that is proportional to absolute temp

27

What’s a barometer

Measures atmospheric pressure
Hg height based on balance forces: gravity (Hg down pull) and atmospheric pressure (pushed Hg up into tube)

28

What’s a manometer

Measures pressure greater than atmospheric pressure

29

Graham’s law of Diffusion/Effusion

Rate of effusion of gas inversely proportional to square root of molar mass
Rate of effusion:
Sqr root of (MM gas 2/MM gas 1)

30

STP

Temp:
O C or 273 K
Pressure:
1 atm (760 mmHg, 760 torr)

31

Molar volume

(nRT/P)
At STO 1 mole of gas occupies 22.4 L

32

Combined gas law

P1V1/T1=P2V2/T2

33

Ideal gas law (including density)

PV=nRT
(T must be made Kelvin)

m/V=(PM/RT)

34

Avagadro’s law

V1/n1= V2/n2

35

Gay-Lussac’s Law

P1/T1=P2/T2
(Temp must be made Kelvin)

36

Charles’s Law

V1/T1=V2/T2

37

Boyle’s Law

P1V1=P2V2

38

Gas density equation

M=(dRT/P)

39

Dalton’s Law of Partial Pressures

Pf=P1+P2+P3...
•is pressure of each has
•pressure depends on #of gas particles but not type
•total pressure exerted by gases in mixture is sum of partial pressure of gases

40

How are real gases different from ideal gases

At high temp: pressure is smaller than predicted; volume is larger
At low temp: pressure is smaller than predicted
Ideal gas behavior has to be corrected when at high pressure/low temp

41

Corrections to ideal gas law (Van Der Waals equation)

(P+(n^2a/V^2))(V-nb)=nRT

42

Define thermodynamics

Study of energy and its transformation

43

Define system

Region of concern

44

Define surroundings

The rest of the universe

45

Define thermodynamic equilibrium

Temperature is uniform throughout material

46

Define enthalpy

Energy associated with breaking/forming bond in chemical reaction
• total energy for constant pressure system
• total energy is internal energy (E) and energy expanded to push surroundings aside
• /\H=/\E+P/\V

47

Define state function

Function dependent on initial/final stages but not pathway

48

Define heat

Energy transferred from higher temp object to lower temp object

49

Define endothermic

Heat flowing into system from surroundings

50

Define exothermic

Heat flows out of system into surroundings

51

Define calorimeter

Used to measure absorption or release of heat by physical change/chemical process
•-qsystem=qcalorimeter

52

Define calorimetry

Experimental measurement of heat transferred during physical or chemical change

53

1st law of thermodynamics

Total energy of universe is constant
•energy can be transferred from system to surroundings or vice vers
•PE—>KE
•Chemical energy—>heat

54

Law of Energy Conservation

Energy cannot be created or destroyed

55

Explain open system, closed system, and isolated system

Open: energy and mass transferred
Closed: only energy is transferred
Isolated: no mass or energy transferred