1. Matter Notes
Matter
and
Change

2. Chemistry, it’s all about matter!
Matter is anything that has volume and mass.
Examples of Matter:
Earth, air, water, books, people, etc.
NOT MATTER:
Light, heat, radio waves, magnetic fields

3. Properties of Matter Focuses on the behavior and composition of matter
Mass
Color
State (solid, liquid, gas)
Ability to dissolve in water
Conducts electricity

4. The Macroscopic View of Matter
Matter that is large enough to be seen
This is the beginning for all of our observations.
We can get hints of what matter is made up of by these types of observations …..
but we need to look closer (smaller) to see the real structure and behavior

5. The Submicroscopic View
You can’t even see this with the most powerful microscopes.
ATOMS!
This is where the majority of our study of chemistry is!
This is why we use …

6. Models in Chemistry
Because we can’t see the submicroscopic matter, we use models as a representation.
There are models of what we think atoms look like and of how we think they behave.
Scientific models are like model airplanes or cars, but these models have been tested and experimentally verified.

7. The Composition of Matter
Qualitative – uses words of descriptions
Color
Heat
Texture
Quantitative – uses numbers and data
Temperature
Length
Mass
Volume

8. Types of Matter Pure Substances – are made of only one type of matter
Elements
Sodium
Potassium
Chlorine
Compounds
Sugar
Salt
Water

9. Mixtures
A physical blend of two or more components that are NOT in a fixed proportion
Examples:
Iced tea
Salt water
Pizza
Salad

10. Pure Substances Compounds Elements
Two or more elements chemically combined
Has different properties than the individual elements
NaCl – table salt
Elements
The building blocks of matter
The simplest form of matter
90 elements are naturally occurring
Fewer than half the elements are abundant enough to play a significant role in chemistry.

11. Elements are organized on the periodic table.
Element symbols are used to represent elements.
Ag: Silver
H: Hydrogen
Formulas are combinations of chemical symbols and their amounts, used to represent compounds.
H2O: Water
Fe2O3: Iron (III) Oxide (rust)

12. Types of Mixtures
Heterogeneous – composition is NOT uniform (pizza, salad, water and oil)
Homogeneous – composition is uniform (Kool-Aid, motor oil)

13. Homogeneous Solutions
Gas/Gas – air (oxygen and nitrogen)
Liquid/Gas – soda (water and carbon dioxide)
Liquid/Liquid – vinegar (water and acetic acid)
Liquid/Solid – salt water (water and salt)
Solid/Liquid – sponge (sponge and water)
Solid/Solid – stainless steel (iron, chromium, and nickel)
Alloy – metal solution

14. Parts of a Solution
Solute – the substance being dissolved, present in a smaller quantity
Solvent – substance doing the dissolving, present in a larger quantity

15. Aqueous Solution
A solution in which the solvent is water

16. Physical Changes
Alters the properties of the material without changing the composition
Phase changes – boil, freeze, melt, condense
Dissolve, cut, grind, bend, break split, crack, crush

17. States of Matter
Solid
Liquid
Gas
Plasma
(ionized gas)

18. Changes in State Freezing – from a liquid to a solid
Melting – from a solid to a liquid
Evaporation – from a liquid to a gas
Condensation – from a gas to a liquid

19. Physical Properties
Properties of materials that can be observed without altering the material.
We use these to separate mixtures.
Solubility, melting point, boiling point, color, density, electrical conductivity, state (solid, liquid, or gas)

20. Volatility
A volatile substance changes from a liquid to a gas at room temperature.
Gasoline
Alcohol

21. Chemical Properties
Properties of materials that can only be observed when the substance changes composition
Cannot be reversed without a chemical reaction
Rusting, Flammability

22. Chemical Changes Alters the composition of the substance
Creates a new substance
Cannot be reversed without another chemical reaction
Burning paper produces soot, carbon dioxide, and water vapor
Combining and acid and a base forms water

23. Summary Matter Mixtures Pure Substances
Hetero. Homo. Elements Compounds
Physical Changes
Chemical Changes

24. Density The ratio of the mass of an object to its volume
Mass/volume (g/mL or g/cm3)
Density of water = 1 g/mL
Less dense objects float, more dense objects sink
Increasing temperature decreases density (hot air rises)

25. Manipulating the Density Equation
D = m/v
m = Dv
v = m/D

26. A bar of silver has a mass of 68. 0 g and a volume of 6. 48 mL
A bar of silver has a mass of 68.0 g and a volume of 6.48 mL. What is the density of silver?
D = m/v
68.0 g / 6.48 mL
D = 10.5 g/mL

27. A copper penny has a mass of 3. 1 g and a volume of 0. 35 mL
A copper penny has a mass of 3.1 g and a volume of 0.35 mL. What is the density of copper?
D = m/v
3.1 g / 0.35 mL
D = 8.9 g/mL

28. Signs of a Reaction: Gas formation (bubbles, fizzing) Color change
Energy (heat/light released or absorbed)
Precipitate formation – solid settling out of a liquid

29. Law of Conservation of Mass
During any chemical reaction, mass is neither created nor destroyed, it is conserved
the mass of the products (what is made from the reaction)
the mass of the reactants (what goes into a reaction)

30. Energy in Reactions Endothermic Exothermic
the capacity to do work; often observed as heat or light
Endothermic
Reaction absorbs heat
Feels cold
Ice packs
Exothermic
Reaction gives off heat
Feels hot
Hand warmers