Transcript Matter notes

Ch 2
Chemistry Basics
Chemistry: The study of Matter
Matter: Anything that has Mass and
occupies Space.
Chemistry Basics
 Mass: the amount of matter an object contains
 Substance: matter that has a uniform and definite
composition.
 Volume: is how much three-dimensional space it
occupies.
Chemistry Basics
 There are both PHYSICAL and CHEMICAL factors
used in chemistry to better understand elements and
compounds found in everyday life.
Physical vs. Chemical
PHYSICAL
 Physical Property: quality or condition of a substance
that can be observed or measured (Ex: color, solubility,
odor, hardness, density, melting point, and boiling
point).
 Physical Change: a change alters the material, but not
its composition (Ex: cutting, bending, and grinding).
Physical vs. Chemical
CHEMICAL
 Chemical Property: the ability of a substance to
undergo a chemical reaction and form new substances.
 Chemical Change: a change alters the material and its
composition (Ex: burn, rot, corrode, rust, explode)
States of Matter
 Matter can come in the form of:
Solid: matter that has a definite shape and volume.
Liquid: a form of matter that flows, has fixed
volume, and takes the shape of its container.
Gas: matter that has no definite shape or volume,
it adopts the shape and volume of its container
LOOK at the TABLE!
A lot of space between
particles. Particles rotate,
vibrate and move around
Little space between
particles. Particles rotate
and vibrate.
No space between
particles. Particles only
vibrate without changing
position.
Chemical Reactions
 Chemical Reaction: one or more substances change
into new substances (Ex: H + O2 make H2O)
 Reactants: starting substances in a reaction
(Ex: H+ O2)
 Products: substance formed from a reaction (Ex: H2O)
Chemical Reactions
Reactants
Products
*Some Chemical Reactions can be reversed, others can not!
**Mass cannot be created or destroyed: Law of Conservation of Mass
Ch 5
Early Atom Study
 Ancient Greeks proposed that matter is made up of
extremely small particles that can not be divided any
further. These particles were called “Atoms” after the
Greek word for “invisible”.
Dalton’s Atomic Theory
 Dalton’s Atomic theory:
a) Atoms of the same element are identical, but
different than other elements.
Vs
Hydrogen (H2)
Oxygen (O2)
b) Compound: combination of different atoms in
whole number ratio.
+
Hydrogen (H)
Ratio of H to O2
=
Oxygen (O2)
2:1
Water (H2O)
Ratios: Examples
NaCl
CH4
1:1
1:4
Structure of the Atom
 J .Thomson used a cathode ray tube to find small
negative particles and named these particles
“electrons”.
 His model was called the “Plum-Pudding Model”,
showing that the rest of the atom not occupied by the
negative charge was positively charged.
Structure Theory
 Lord Ernest Rutherford directed alpha particles (positively
charged particles much smaller than an atom) at a thin
piece of gold foil. Some particles went straight through
while others were slightly deflected (thought to happen if
Plum Pudding Model), yet some were greatly defected or
repelled backward. This showed that there was a central
positive nucleus (positive repels positive).
Rutherford’s Conclusion
 Atom consisted of mostly
empty space.
 Each atom contained a small
dense positively charged
central portion: Nucleus
 This Nucleus was always
positively charged and made
up of protons (+) and
neutrons (no charge).
 Atom was always neutral (no
charge.
Modern Atomic Theory
 Bohr’s Model: electrons move around the nucleus in a
manner similar to that of the planets moving around
the Sun (planetary model).
Modern Atomic Theory
 Wave-Mechanical Model: electrons move around in
areas called orbitals.
 Orbitals: the region in which an electron with a
particular amount of energy is most likely to be
located.
The Periodic Table of Elements
Examples:
Protons: 6
Electrons : 6
Neutrons: 12-6= 6
Nuclear Charge: +6
*Atom is neutral
Protons: 11
Electrons : 11
Neutrons: 23-11= 12
Nuclear Charge: +11
*Atom is neutral
Find the Element:
Protons= 13
Electrons= 13
Neutrons=14
Answer:
Aluminum
Al
With an atomic mass of: 13+14=27
Isotopes
 Isotopes: atoms with the same atomic number (the
same protons and electrons), but different atomic
mass (different neutrons).
Example:
If there is an isotope of Carbon with an atomic mass of
13. How many:
Protons= 6
Electrons= 6
Neutrons=7
**Protons never change!!! Only Electrons (charge) and
Neutrons (atomic mass).
Special Isotope: Hydrogen
 Hydrogen is the only element whose Isotopes
have special names.
 Hydrogen with an atomic mass of 2= Deuterium
 Hydrogen with an atomic mass of 3= Tritium
Question
 Carbon has 2 naturally occurring stable isotopes. Most
carbon atoms (98.89%) are C-12, while the remaining are
C-13 (1.108%). What is the atomic mass of Carbon?
98.89 x 12
= 11.87 amu
100
1.108 x 13 = 0.1440 amu
100
11.87 + 0.1440= 12.01 amu (same as on table!)
Location of Electrons
 Electrons are located in areas called orbitals. Orbitals
form a series of energy levels. Number of the period
shows the number of the energy level.
Outer Electrons
 Valence electrons: are the outermost electrons in an
atom—the ones that are involved in bonding.
Atomic Orbitals
 When electrons occupy the lowest available orbitals, the
atom is in the ground state.
 When electrons absorb energy (gain), it can temporarily
move to a higher energy level and go into an excited state.
 When electrons return to ground state, they emit (give off)
the same amount of energy. This is called a spectrum line
or bright line spectrum.
Spectra
 Each atom (element) has its own spectrum and can be used
to identify elements.
Extra Vocabulary to know:
 Element: substances that cannot be broken down or
decomposed into simpler substances by chemical means.
 Atom: the smallest particle of an element that can enter
into a chemical reaction.
 Molecule: combination of more than one atom
 Compound: a substance composed of two or more
elements that are chemically combined in definite
proportion by mass.