1. 2.1 The Nature of Matter

2. Atoms The most basic unit of matter
First proposed by Greek philosopher Democritus nearly 2500 years ago.
Atom comes from the Greek word atomos which means “unable to be cut.”
100 million atoms in a row is 1 centimeter long.
Composed of 3 subatomic particles, protons, neutrons, and electrons.

3. Subatomic Particles Protons In the nucleus Positively charged
# protons = atomic # (identifies the element)
Neutrons
Neutral (no charge)
Mass number = # p + # n
Electrons
Negatively charged
Outside the nucleus
# e- = # p (atom is neutral!)

4. Elements and Isotopes Elements:
Pure substances that contain only one type of atom.
Metals, non metals and metalloids
Isotopes are elements with a different number of neutrons which will change the mass of the element.
Some isotopes are radioactive, so their nuclei break down.
Common elements essential to life:
Oxygen Carbon, Nitrogen, Phosphorus , Calcium, Sodium, Iodine, Magnesium.

5. The Periodic Table

6. Elements and Isotopes Similarities of elements and isotopes
Differences between elements and isotopes
Name
Atomic Number
Protons (change the proton change the element)
Electrons (change the electron change the charge)
Neutrons
Mass
Neutrons are in the nucleus with protons and make up the mass, change the neutral neutrons and you change the mass.

7. Chemical Compounds
Substance formed by the chemical combination of two or more elements in definite proportions.
The properties of compounds are different than the elements from which they are formed

8. Chemical Bonds Electrons in the outermost shell of atoms form bonds
Ionic bonds are between a metal and a non metal – bond formed by transfer electrons; cation (+) and anion (-)
Covalent bond between non-metals – bond formed by the sharing of electrons between the atoms

9. Ionic Bonds
Cations are metals that lose their electrons and become positive. They become positive when they lose negatively charged electrons.
Anions are non metals that gain electrons.
Opposite charges (negative and positive) held together with this transfer of electrons.
Ionic Compounds:
Solid at room temperature
Form crystals
Conduct electricity when dissolved in water
High melting point.

10. Covalent Bonds 2 non metals – electrons shared between adjacent atoms
Polar bonds
Unequal sharing
Slightly negative and slightly positive regions of the molecular which allows for intermolecular forces like Hydrogen bonding. Opposites attract!
Non Polar- these elements share their electrons equally.

12. Types of Bonds

13. Intermolecular forces
Attractive forces between molecules
Hydrogen bonding (strong)
Dipole-dipole interactions
Van der Waals forces (weak)

14. Intermolecular forces
Hydrogen bonding
When H bonded to O, N, or F, H on one molecule attracted to O, N or F on another molecule
Water’s properties due to H-bonding
Holds DNA together

15. Intermolecular forces
Dipole – Dipole Interactions
Attractive forces between dipoles in polar molecules

16. Intermolecular forces
Van Der Waals forces
Weaker interaction
Differences in electron density around the molecule create slightly positive/negative regions
All covalent compounds exhibit van der Waals forces
Allows lizards to walk on walls.

17. Intermolecular forces
When molecules are close together, and electrons are not shared equally, a slight attraction can develop between the oppositely charged parts of the molecules. Opposites attract!
This molecular polarity allows for molecules to stick together!