2. • Chemistry – study of all forms of matter
(solids, liquids, gases, plasma)
• Physics – study of the way energy affects
matter
• Property is the ability to do something
• Can be observed or measured without
changing the matter’s identity
3. Matter
• Matter is made of atoms in motion
• Matter is anything that has mass and takes up
space (volume)
• 4 states of matter – solid, liquid, gas, plasma
• Water is the only compound that exists on earth
in all 3 states
• No 2 objects can occupy the same space at the
same time
• Inertia is a property of matter: moving doesn’t
want to stop, stopped doesn’t want to move
4. States of Matter
• Solid: either crystalline or amorphous, can not
change shape, has a fixed volume
• Liquid: particles slide past each other, change
shape, not volume
• Surface tension and viscosity are properties
of a liquid
• Gas: particles are spread out, change shape
and volume, a gas is the only state of matter
that can be compressed
5. Density/solubility
• Density is mass/volume and identifies how
tightly packed matter is.
• Water is the only substance with a density of
1g/mL.
• Density is unique to an element or a
compound.
• Solubility- the ability to dissolve at a given
temperature and pressure
6. Examples of physical properties and
physical changes
• Physical Property
• Boiling Point
• Freezing Point
• Melting Point
• Condensation Point
• Evaporation Point
• Ductile
• Malleable
• Solubility
• electric conductivity
• density
• Physical Change
• Boiling
• Freezing
• Melting
• Condensing
• Evaporating
• Stretching into wire
• Hammer into sheet
• Dissolving
• Moving electricity
• Mass/volume
7. NOT physical properties
• Never use weight, mass, volume, shape or
size to describe matter
• Weight is a measure of the force of gravity,
relative to location
• In other words, physical changes do not
change the identity of the matter
8. Chemical Properties
• Chemical property – a property of matter that
describes a substances ability to participate in
chemical reactions (you get something new)
• Examples of chemical properties
– Flammability – ability to burn
– Reactivity with acid
– Reactivity with water
– Reactivity with oxygen
– Reactivity with other elements
– Combustion (ignite, or burst into flame)
9. Evidence of Chemical Change
• Heat produced
• Cold produced
• Gas produced
• Light produced
• Odor change
• Color change
• Precipitate produced (solid)
• Oxidation (tarnish)
• Electrolysis
• Bubbling
• Foaming
• Fizzing
10. Mixtures, Solutions, Suspension, Elements,
Compounds, Acids, Bases and Salts
• Mixture: Two or more substances that are not
chemically combined
• All mixtures can be physically separated
• Ratio of mixtures are not fixed
• Mixtures can be solid, liquid or gas
• Solution: Mixture that appears to be a single
substance
• Material must be soluble (able to dissolve)
11. • Solute is what is dissolved
• Solvent what the solute is dissolved in
• Water is the universal solvent
• Solubility is the ability of substances to
dissolve at a given temperature and pressure
• Suspensions: are mixtures where the particles
are heavy enough to settle out (sink to
bottom) of the solution, scatter light, can be
filtered
12. Elements
• Elements: are pure substance that cannot be
separated into simpler substances by physical or
chemical means
• Elements can be identified by their characteristic
properties (physical and chemical)
• Elements are classified by large categories:
• Metals – shiny, good conductors
• Nonmetals – dull, poor conductors
• Metalloids –has properties of metals and
nonmetals depending on conditions
13. Compounds
• Compounds: pure substance composed of two or
more elements that are chemically combined
• Elements do not form compounds randomly
• Compounds form in specific mass ratio
• When elements form compounds, new
characteristics properties are created
• The only way to separate a compound into
elements or other compounds is by a chemical
reaction which allows for a chemical change by
adding (endothermic) or taking away
(exothermic) energy
14. Acids, Bases and Salts
• Most acids start with the element H, hydrogen
– EX: HCl (hydrochloric acid), H2SO4 (sulfuric acid)
• Any compound that increases the number of
hydronium ions, (H+), when dissolved in water is an
acid
• NEVER TASTE, SMELL OR TOUCH ACIDS
• Properties: has sour taste (think vinegar or
lemons), corrosive, react with some metals to
produce hydrogen gas, conduct electricity
• Most widely made acid H2SO4
15. • Most bases end with OH-, a hydroxide ion, when
dissolved in water
• NEVER TASTE, SMELL OR TOUCH BASE
• Properties: has bitter taste and slippery feel
(think soap), corrosive, conducts electric current
• Strong acids will have additional Hydrogen (H+)
molecules in the compound (HCl – vs- H2SO4)
• Strong bases will have additional hydroxide (OH-),
molecules in the compound (MgOH –vs- Ba(OH)2)
• 7 on the pH scale is neutral (H2O)
• Bases have a pH greater than 7
• Acids have a pH less than 7.
16. Salts
• Large group of compounds with similar
properties (usually formed with elements in
group 17, halogens)
• When a reaction occurs between an acid and
a base, they neutralize each other
• When an acid neutralizes a base, salt and
water are produced
17. Atomic Theory
• Democritus (440 BCE)- realized that if you
continued to cut something, eventually you
would end up with something that couldn’t be
cut anymore, atomos – meaning not able to
divide
• Atoms are the smallest particle that an
element can be divided and still be the same
substance
• All matter is made of atoms
18. John Dalton (1803)
• realized that atoms combine in very specific
proportions (ratios) based on mass
• all substances are made of atoms and they can not
be created, divided or destroyed because they were
made of a single substance
• All atoms of the same element are exactly alike and
different from other elements, they are unique
• Atoms join with other atoms to form new substance
19.
20. J. J. Thomson (1897)
• discovered that there were small particles
inside the atom, meaning that atoms can be
divided into smaller substances
• Electrons – negatively charged particles
attracted to positively charged particles
• Plum pudding model – electrons are mixed
throughout the atom, soft blobs of matter
21.
22. Ernest Rutherford (1909)
• Discovered that an atom contains a nucleus
with positively charged particles and that the
electrons must be “floating” around the
nucleus
• Most of an atom is empty space
23.
24. Niels Bohr (1913)
• Proposed that electron moved around the
nucleus in energy levels (shells), but no
electrons between the energy level (think
ladder)
• Electrons can jump from one level to another
• Travel in a definite path
25.
26. Modern Atomic Theory
• Erwin Shrodinger & Werner Heisenberg
• Electrons have no predictable pattern and
move in a region where electrons are likely to
be found called the electron cloud
27.
28. Atoms
• All atoms have a nucleus
– protons (+),
– neutrons (no chg)
– electrons (-)
• Same number of protons and electrons an atom has
no charge
• More protons (+) than electrons (-) the atom has a
positive ion is formed (more positives than negatives)
• More electrons (-) than protons (+) a negative ion is
formed (more negatives than positives)
29. • 117 different element that are unique and all
things known to exist come from a
combination of these elements in specific
mass ratios
• Simplest atom is made of one proton, and 1
electron – hydrogen (has no neutrons)
30. • All additional element will have protons,
neutrons and electrons
• The atomic number of an element is
determined by the number of protons,
– 1 is hydrogen, 6 is carbon, hydrogen has 1 proton,
carbon has 6 protons (you can not change the
number of protons)
• To find neutrons take the mass number
(rounded) and subtract the protons.
31. Isotopes
• Isotopes have the same number of protons but
additional neutrons which causes the atomic mass to
be different
• Isotopes can be stable (maintain there structure) and
unstable (fall apart over time)
• Unstable isotopes are radioactive and will decay over
time giving off particles and energy (radioactive)
32. • Mass number determines the isotope, the
number of protons and neutrons added
together
• Most elements have isotopes
• All isotopes of an element have the exact
properties of the element
33. Forces in atoms
• Gravitational force – pulls objects toward each
other—depends on mass and distances
between the objects—very small force in
atoms
• Electromagnetic force –– proton (+) and
electrons (-) have strong attraction which
keeps the electrons in motion around the
nucleus of atoms
34. • Strong force – force which keeps protons from
flying apart due to close distance between
protons and neutrons
• Weak force – relevant to radioactive atoms-allows
neutrons to change into proton and
electron
35. Periodic Table
• Dmitri Mendeleev-recognized that elements
had repeating patterns (periodic) and
organized elements into a table by increasing
atomic mass
• Henry Moseley - determined that the number
of protons - atomic number (which is unique
to each element) would allow the elements to
fit into very specific pattern
36. • Separated into 3 large categories: metals,
metalloids, nonmetals based on their
properties, moving from left (very reactive) to
right (gradually becoming completely non-reactive)
across each period on the table
• Columns are called groups or family, each has
a name
• Each element in a family has the same
number of valence electrons in the outer shell
37. • Group 1 – Alkali Metals
• Group 2- Alkaline Earth Metals
• Group 3-12 – Transition Metals
• Group 13 – Boron Group
• Group 14 – Carbon Group
• Group 15 – Nitrogen Group
• Group 16 – Oxygen Group
• Group 17 – Halogens
• Group 18 – Noble Gases
• 1st Row at bottom – Lanthanides
• 2nd Row at bottom - Actinides
38. • Rows (left to right) are called periods (7 rows)-
determines the number of energy levels
• 1st energy level – 2 valence electrons (max)
• 2nd energy level – 8 valence electrons (max)
• 3rd energy level – 18 valence electrons (max)
• And so on….
• Each energy level can have less valence
electrons but they can not have more than
the maximum valence electrons.
39. Bonds – Octet Rule
• To form bonds, elements must reach a full
state of 8 valence electrons in the outermost
energy level (octet rule) (Exception: would be
first energy level which is full at 2-helium)
40. • Atomic number = Number of Protons
• Electrons equal to the number of protons
• Neutrons equal atomic mass (rounded) minus
the protons
• Protons do not change in a atom, neutrons
can change (isotopes), electrons can be shared
or transferred (when bonds are made)
41. Chemical Bonding
• Bonds are formed between elements that can
rearrange their molecules to create compounds.
• Bonds only form when energy is put in or taken
away (endothermic or exothermic)
• For bonds to form, the outer energy level must
equal 8 (octet rule) for one of the elements.
• Ionic bond – electrons are transferred (one
element gets to 8) between a metal and a
nonmetal
• Covalent Bond – electrons are shared (both
elements get to 8) both nonmetals
42. • If any atom losses electrons, the molecule
becomes positively charged ion.
• If an atom gains electrons, the molecule
becomes negatively charged ion.
43. Steps to Draw Bonds
• Determine the number of valence electrons
• Determine which elements needs what
• Draw the electron dot diagram with the electrons
to be transferred or shared in the middle
• Draw the bond – Vin diagram
• Determine if the bond is ionic (one reaches 8)
metal bonded to non-metal or covalent (both
reach 8), non-metal bonded with non-metal
• Determine whether a single, double or triple
bond has occurred
44. Steps to Draw Molecular Structures
• Determine the elements
• Determine the atoms
• Determine the number of molecules
• First elements is always the center of the
structure
• Additional elements are drawn around the
center
