Chemistry

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1

Homogeneous Mixture

Mixture with uniform density throughout and no distinguishable coponents.

2

Heterogeneous Mixture

Mixture in which the components are readily distinguished.

3

Physical Change

A change in which the checmical composition of a substance remains the same.

4

Chemical Change

A change in which the chemical bonds are broken and reformed to create a new and different substance.

5

Element

The simplest of substances and is represented by a specific letter or combinations of letters.

6

Compounds

Combinations of eleements in whole number ratios.

7

Law of Conservation of Mass

Mass cannot be created or destroyed during a chemical reaction.

8

Chemical Reaction

The breaking of bonds and the reforming of new bonds to create new chemical compounds with different chemical forumulas and different chemical properties.

9

Five main types of chemical reactions.

syntehesis, decomposition, combusion, single replacement, double replacement

10

Syntehsis Reaction

Two elements combine to form a product.

11

Decomposition Reaction

The breaking of a compound into component parts.

12

Combustion Reaction

The reaction of a compound or element with oxygen. In the combusiton of a hydrocarbon carbon dioxide and water are produced.

13

Replacement Reactions

Reaction involoving ionic compounds. The reactivity of the ionic compounds determines whether the reaction will take place or not. Can be single replacement or double replacement reaction.

14

Single Replacement Reaction

Reaction between a more active metal reacting with an ionic compound containing a less active metal to produce a new compound ex. copper wire reacting with aqueous silver nitrate.

15

Double Replacement Reaction

Reaction involving two ionic compounds where the positive ion from one compound combines with the negative ion of the other compound. The result it two new ionic compounds that have switches partners.

16

Atomic Number

The number of protons in a given element.

17

Atomic Weight/ Atomic mass number

An average of the masses of each of the iostopes of an element as they occur in mature. (Represents the number of protons and the number of neutrons in an element because electrons essentially have no mass).

18

Calculating the number of neutrons in a given isotope of an element.

Subtract the atomic number from the mass number.

19

Charge of Elements in Group IA

+1

20

Charge of Elements in Group IIA

+2

21

Charge of Elements in Group IIIA

+3

22

How do you calculate density?

d = m/v

23

Specific Gravity

Density of an object realtive to water. No units.

24

Kinetic Energy

Energy of motion.

25

Potential Energy

Stored energy.

26

Heat

Form of energy, measured in calories.

27

calorie

The amount of heat necessary to raise the temperature of 1g of water by 1 degree celsius.

28

Mixture

Combination of two or more pure substances.

29

Isotope

Atoms with the same number of protons but different numbers of neutrons.

30

Noble Gases

Elements in group 8A of the periodic table. Have no charge and are gases under normal conditions. (Helium, Neon, Argon, Krypton, Xenon, Radon)

31

Electron Configuration

Shell Electrons (max)
1st = 2
2nd = 8
3rd = 18
4th = 32

32

Halogens

Elements of group 7A. Have a charge of -1. Fluorine, Chlorine, Bromine, Iodine. Form compounds with sodium in the form NaX.

33

Alkali Metals

Elements of group 1A. Have a charge of +1. Lithium, Sodium, Potassium, Rubidium, Cesium. React with water to form hydrogen gas and a metal hydroxide (MOH) + H2. Also from compounds with the halogens in the form MX. Ex. NaCl.

34

Anion

An ion with a negative charge.

35

Cation

An ion with a positive charge.

36

Ionic bond

A chemical bond resulting from the attraction between a positive ion and a negative ion.

37

Covalent Bond

A chemical bond resulting from the sharing of electrons between two atoms.

38

Polar covlaent bond

A covalent bond between two atoms where electrons are not shared equally between the two atoms.

39

Dipole

Created when atoms are joined by a polar covalent bond. The positive end of a dipole in one compound will be attracted to the negative dipole in another compound creating weak attraction between the two compounds.

40

Strongest type of chemical bond.

Covalent bond

41

Weakest typ of chemical bond.

Ionic bond

42

Strongest of intermolecular forces.

Hydrogen bond

43

Weakest of intermolecular forces.

Dispersion forces.

44

Hydrogen bond.

Attraction for a hydrogen atom by a highly electronegative element. Generally involve fluorine, chlorine, oxygen, and nitrogen.

45

Dispersion Forces

Temporary dipole created when moving electrons within an element or compound concentrate themselves on one side of an atom. Usually found in nonpolar covalent compounds.

46

Van der Walls forces

Another name for dispersion forces, dipole interactions.

47

Charge of elements in group VA.

-3

48

Charge of elements in group VI A

-2

49

Charge of elements in group VII A

-1

50

How many known elements are there?

109

51

What is a mole.

An amount of an element equal to its atomic weight in grams. Also described by the amount of a substance that contains 6.02 x 10 23rd particles of that substance.

52

What is alpha radiation?

The emission of helium ions that consist of 2 protons and 2 neutrons (thus having a +2 charge). Alpha particles can be stopped by a piece of paper.

53

What is beta radiation?

The product of the decomposition of a neutron and is composed of high energy high-speed electrons. They are negatively charged and have basically no mass. Beta particles can be stopped by aluminum foil.

54

What is gamma radiation?

High-energy electromagnetic radition that lacks charge and mass. Gamma radiation can be stopped by several feet of concrete or several inches of lead.

55

Radioactivity

The emission of particle sfrom an unstable nucleus. Exists in three forms alpha, beta, and gama radiation.

56

Molar mass

The mass of one mole of a compound.

57

Sublimation

When a substance changes from a solid to a gas without first becoming a liquid.

58

Ideal gas law

PV=nRT (n is equal to the number of moles of the substance and R is the gas constant 0.082)

59

Molarity

The number of moles of solute in 1 liter of solution.

60

Redox reactions

Reactions that involve the transfer of electrons from one element to another.

61

Oxidation

The loss of electrons in an redox reaction.

62

Reduction

The gain of electrons in a redox reaction.

63

Rules for determining oxidation state.

1. Elemental atoms have an oxidation number of zero.
2. The oxidation number of any simple ion is the charge of the ion.
3. The oxidation number for oxygen in compound is always -2.
4. The oxidation number for hydrogen in compound is +1.
5. The sum of the oxidation numbers equals the charge on the molecule or polyatomic ions.

64

Acids

- act as hydrogen-ion donors.
- produce H3O+ in aqueous solutions.
- tast sour or tart.
- most of their formulas begin with H.
- relase H2 gas when reacting with active metals.
- conduct electrical current.
- pH is less than 7.

65

Bases

- produce OH- in solution.
- taste bitter.
- feel slippery.
-conduct electricity.
- formulas often contain OH-.
- pH is greater than 7

66

Neutralization

Process which occurs when an acid and a base react tot produce a salt and water. The result is a pH near 7.