1. CHEMISTRY: The Study of Matter

2. Describing Matter Chemistry: study of matter
Matter: has mass and takes up space
Types of matter:
pure substances (sugar, salt)
not pure substances (milk, fruit)

3. States of Matter: GASES
Volume can change easily
Gas particles spread apart filling all of the space available
Do not have a definite shape or a definite volume

4. GAS BEHAVIOR Boyle’s Law:
Increased pressure= decreased volume. Decreased pressure= increased volume.

5. GAS BEHAVIOR Charles’ Law: Increased temperature = increased volume
Decreased temperature = decreased volume

6. IN SUMMARY….
Gases at high temperatures have high pressure and a larger volume
Gases at low temperatures have low pressure and low volumes.

7. States of Matter: SOLIDS AND LIQUIDS
Solids: definite shape, definite volume.
Liquids: definite volume, but no shape
solids
liquids

8. States of Matter: SOLIDS AND LIQUIDS
Like gases, solids and liquids expand when heated and shrink when cooled.
What evidence do you have of this?
Can you think of an exception?

9. Changing State deposition melting evaporating condensing freezing
sublimation

10. Describing Matter Atoms: The most basic particles that make up matter
Elements: pure substances
Cannot be broken down into
smaller pieces
protons
neutrons

11. Describing Matter
Compounds: two or more elements chemically combined (NaCl,sodium chloride, table salt)
Mixtures: two or more substances, not in a specific pattern (salad dressing, dirt).
can be homogeneous (same throughout- salt water) or heterogeneous (not the same throughout- salad dressing)

12. Measuring Matter Weight: how gravity affects the matter in an object
Mass: the amount of matter in an object, not affected by gravity (unit= gram)
Volume: the amount of space matter occupies (unit = milliliter, liter)
Density: how tightly matter is packed into a given space (unit = gram / ml or l)
Density = Mass / Volume

13. Describing Matter Two Properties of Matter:
Physical Properties: observed without the matter changing (state of matter, color, texture, flexibility, solubility)
Chemical Properties: observed when matter changes into new substances (flammability, reactivity)

14. Changes in Matter
Physical Changes: changes the form or appearance of matter, but does not make the matter into a new substance (ice melting)
Phase changes are physical changes
deposition
melting
evaporating
condensing
freezing
sublimation

15. Changes in Matter
Chemical Changes: Produces new substances (candle burning).

16. Density Density relates the mass of a material in a given volume.
Density = Mass ÷ Volume (g/mL)
Objects with a HIGH density will sink in water
Objects with a LOW density will float in water

17. Atoms and the Periodic Table
Atom: smallest particle of an element
Have three, subatomic particles: protons (+), neutrons (), and electrons (-).
Protons and neutrons: in the center of an atom (nucleus) surrounded by electrons in a cloud.

18. Scale and Size of Atoms Amazingly small
Tiniest speck of dust contains 10 million billion atoms

19. Atomic Number Every atom of an element has the same number of protons.
Unique number of protons = atomic number.
Atomic number identifies an element. Carbon (C) = 6, Oxygen (O) = 8, Iron (Fe) = 26.
If an atom has to be stable, how can you figure out the number of electrons?

20. Organization of the Periodic Table
Properties of an element can be predicted from its location on the PT
Periods: Horizontal rows on the PT (elements lose reactivity as you move across periodic table, until you reach the gases)
Groups: Vertical columns on the PT (elements in groups have similar properties)

22. Reading an Element’s Square

23. Metals
Physical Properties: shiny, malleable (bendable), ductile (stretchy), conductive (transfers heat or electricity)
Chemical Properties: react by losing electrons to other atoms (the fewer electrons to lose or gain, the more reactive an element is)

24. Metals
Some are extremely reactive (sodium) others are not reactive at all (gold, platinum)
Some have moderate reactivity (iron- corrosive)
Reactivity decreases as you move from left to right across the periodic table

25. Alkali Metals
Group 1: react with other elements by losing one electron.
-So reactive, never found as uncombined elements in nature
-Two most important: Sodium and potassium

26. Alkaline Earth Metals Group 2 React by losing two electrons
Two most common are magnesium
and calcium
Each is fairly hard, gray-white, and a
good conductor

27. Transition Metals Metals in Groups 3-12
Most of familiar metals (iron, platinum, copper, nickel, silver, gold)
Most are hard and shiny, good conductors
Less reactive as you move to the right

28. Misc. Metals Some elements in groups 13 and 14 (aluminum, tin, lead)
Not very reactive
Lanthanides (top row in rows below PT): soft, malleable, shiny, highly conductive).
Actinides (row below Lanthanides): Only Actinium (Ac), Thorium (Th), Protactinium (Pa) and Uranium (U) occur naturally

29. Nonmetals Nonmetal: lacks most of the properties of a metal (not
conductive, reactive, dull, brittle)
Carbon Family (group 14):
Elements can gain, lose, or
share four electrons
Very versatile

30. Nonmetals Nitrogen Family (group 15): Gain 3 electrons
Occur in nature in the form of
diatomic molecules
(Di = two, atomic = of atoms).
Ex: P2,N2,

31. Nonmetals Oxygen Family (group 16): Gains or shares 2 electrons
Because oxygen is highly reactive,
it can combine with most other
elements in nature
Oxygen you breathe is diatomic
O2, or can be triatomic, O3, ozone.

32. Nonmetals
Halogen Family (group 17): Halogen means salt-forming (bonds to metals to form salts)
Gains 1 electron
Very reactive, dangerous (but can bond to other elements to be useful)

33. Nonmetals Noble Gases (group 18): Stable
Exist in the Earth’s atmosphere in small amounts
Don’t bond to other elements (have
enough electrons)

34. Hydrogen Simplest element with the smallest atoms
Unique properties, cannot be grouped with other gases
Makes up 1% of Earth’s crust
Rarely found on Earth as pure element
Mostly found as H20 (water)

35. Metalloids Characteristics of metals and non-metals
Most common is Silicon (Si)
Most useful property is their varying ability to conduct electricity
Semiconductors (used in computer chips) are made of metalloids

36. Isotopes
Isotopes: identified by the mass number of an element (the number of protons plus neutrons)
Carbon is most commonly C-12, but can be C-13 and C-14. What’s the difference?

37. Isotopes and Mass Number
Atoms of an element always have the same number of protons, but can have different numbers of neutrons.
Atoms of an element with different numbers of neutrons: isotopes

38. Phase Changes and Energy
The phase a substance depends on the kinetic energy of the molecules
Substances change from phase to phase in response to the input or output of energy
Phase changes reflect a change in energy or a change in the behavior of the atoms or molecules, not a change on the atoms or molecules themselves.
Ex: ice melting, water condensing

39. Phase Changes and Energy
Matter can change from one state to another when thermal energy is absorbed or released.

40. Law of Conservation of Energy
During any chemical or physical change, energy cannot be created or destroyed.
Energy can change forms and can be transferred from one object to another, but the overall amount must be conserved.

41. Temperature
A measure of the average kinetic energy of individual particles of matter
Measured using a thermometer (heat increases the kinetic energy of the liquid inside, takes up more space).

42. Thermal Energy and Heat
The total energy of all of the particles of an object is called Thermal Energy.
The more particles an object has at a given temp, the greater the thermal energy.
The transfer of thermal energy from matter at a higher temperature to matter at a lower temperature is called Heat.
Why does ice melt in your hands?

43. Heat Transfer Heat is transferred in 3 ways:
Conduction: Heat is transferred without the movement of the matter (metal handle of a pot getting hot).
Convection: Heat is transferred by the movement of currents within a fluid (plate tectonics).
Radiation: The transfer of energy by electromagnetic waves. Radiation does not require matter to transfer energy.

44. Examples of Heat Transfer

45. Direction of Heat Transfer
Heat will flow from the warmer object to the colder object
Heat transfer occurs in only one direction
In summary, heat always flows from higher energy levels to lower energy levels.

46. Phase Changes and Energy
Energy transferred TO the substance
deposition
melting
evaporating
freezing
condensing
sublimation
Energy transferred FROM the substance

47. Calories Heat is measured in Calories.
1 Calorie is the amount of heat needed to raise the temperature of 1gram of water 1 degree Celsius.
Equilibrium: when there is no net energy transfer within a system

48. Heat of Fusion
Water molecules in ice are held together by forces called “bonds.”
Energy is required to break the bonds when changing solid water (ice) to liquid water.
The energy needed to change solid water to its liquid form is called Heat of Fusion.
The Heat of Fusion for water = 80c/g (calories per gram)

49. Observing Chemical Change
Matter has chemical and physical properties and undergoes chemical changes and physical changes.

50. Chemical Changes Occur When…
Chemical changes occur when bonds break to form new bonds
Chemical changes occur through chemical reactions

51. Evidence for Chemical Reactions
Two main changes occur during a chemical reaction:
1) Formation of new substances
2) Changes in energy

52. Evidence for Formation of New Substances
Color Change
Precipitate Formed (solid)
Gas Produced

53. Evidence for Changes in Energy
Heat Produced (reactions can be exothermic or endothermic)

54. Matter and Energy
When a reaction causes energy to be released (often as heat), the reaction is called exothermic. When energy is taken in, a reaction is called endothermic (ice melting).
Endothermic
Energy taken in, resulting temp is higher
Exothermic
Energy released, resulting temp is lower

55. Describing Chemical Reactions
Chemical Equations: A short, easy way to show a chemical reaction, using symbols instead of words.
Structure of Chemical Equations: All chemical equations have a common structure. Substances at the beginning are called “reactants” and substances at the end are called “products”

56. Chemical Equations reactant + reactant  product + product
You read the arrow () as “yields.”

57. Law of Conservation of Mass
Chemical Equations must be balanced because of the Law of Conservation of Mass (matter can neither be created nor destroyed)

58. Steps for Balancing Equations
1) Write the equation
H2 + O2  H2O
reactants products
Count the atoms
H2 = 2 hydrogen atoms (reactant)
O2 = 2 oxygen atoms (reactant)
H2 = 2 hydrogen atoms (product)
O = 1 oxygen atom (product)

59. Steps for Balancing Equations
Use coefficients to balance atoms (a coefficient is a number placed in front of a chemical formula to balance products and reactants)
H2 + O2  2H2O
**Coefficients can be placed on the reactant side or the product side of an equation
4) Look back and check

60. Checking your Equations

61. Chemical Reactions and Bonding
Bonds are broken and formed during chemical reactions.
If an atom loses or gains an electron it is called an ion (an element with a + or – charge).
Bonds that form between ions are called ionic bonds (usually between metals and nonmetals).

62. Ionic Bonding

63. Chemical Reactions and Bonding
If elements bond by sharing electrons, covalent bonds form.
Covalent bonds usually form between nonmetals.

64. Covalent Bonds

65. Solutions Solutions: a well-mixed mixture.
Solvent: What does the dissolving (the larger part of a solution).
Solute: What is dissolved (the smaller part of a solution).
Liquid water solution

66. Solutions The effects of solutes on solutions:
Lower freezing points, Higher boiling points.
What affects solubility?
Temperature, pressure, type of solute and solvent

67. Concentration of a Solution
Saturated Solution: A solution that contains as much solute dissolved as possible. (No more can be dissolved)
Unsaturated Solution: A solution where more solute can still be dissolved.

68. Acids and Bases
Properties of Acids: tastes sour, reacts with metals and carbonates, and turns litmus or pH paper red.
Uses of Acids: vitamins,
batteries (sulfuric acid),
fertilizers

69. Acids and Bases
Properties of Bases: taste bitter, slippery, turns indicator paper blue.
Opposite of acids: Do NOT react with metals and carbonates
Uses of Bases:
Cleaning Products

70. Acids and Bases The pH scale used to determine
the strength of an acid or a base.
The pH scale measures the
concentration of hydrogen ions
in a solution.
A low pH: very acidic.
A neutral solution (water) pH = 7
A high pH: very basic.

71. Acid-Base Reactions
When an acid and base react, the product is neither acidic or basic.
Neutralization occurs when an acid and base react (pH becomes closer to neutral)
Salts form as a product of an acid base reaction when the positive ion in a base bonds with a negative ion of an acid.
Ex: HCl + NaOH  H2O + (Na+ + Cl-)