Download PP atoms - Lake County Schools

Chemistry
Chapter Three – Atoms: The
Building Blocks of Matter
South Lake High School
Science Department
Ms. Sanders
Foundation of Chemistry
• Law of Conservation of Mass – mass is neither created nor destroyed
during ordinary chemical reactions or physical changes
• Law of Definite Proportions – chemical compounds contain the same
elements in exactly the same proportions by mass regardless of
sample size or source of compound
• Law of Multiple Proportions – if two or more different compounds
are composed of the same two elements, then the ratio of the
elements is always a ratio of small whole numbers
Foundation

Law of Conservation of Mass
Foundation

Law of Multiple Proportions
Atomic Theory
 Dalton’s Atomic Theory
1. All matter is composed of atoms
2. Atoms of the same element are identical whereas atoms of different
elements differ
3. Atoms cannot be subdivided, created, or destroyed
4. Atoms of different elements combine in simple whole-number
ratios to form chemical compounds
5. In chemical reactions, atoms are combined, separated, or
rearranged
 Modern Atomic Theory
 Atoms are divisible into smaller particles (quarks, mesons)
 One element can have atoms with different masses (isotopes)
Atom Structure
 Atom – smallest particle of an element that retains the
chemical properties of that element
 2 Regions
1. Nucleus – make up the mass of the atom
 Positive protons
 Neutral neutrons
2. Area surrounding Nucleus
 Negative Electrons
 Subatomic particles = protons, neutrons, and electrons
Discoveries
 Discovery of electron was a result of electricity and matter
investigations
 Cathode – ray tubes (cathode - negative, anode – positive)
particles that compose cathode rays are negatively charged
(electrons)
 J.J. Thomson – concluded all cathode rays are composed of
identical negatively charged particles – electrons
Discoveries

Cathode Ray
Discoveries
 Robert Millikan measured the charge of an electron, led to ability to
calculate the mass of an electron
 Ernest Rutherford (associates Geiger and Marsden) bombarded a thin
piece of gold foil with fast-moving alpha particles, some particles deflected
therefore they concluded that there is a small space with a positive electric
charge within atoms – atom mostly empty space with a small, positive
center
Discoveries

Gold Foil Experiment
Discoveries

Gold Foil Experiment
Counting Atoms
 Atomic number – number of protons of each atom of that element
 Found on the periodic table
 Atomic # = # Protons
 Atomic # = # Electrons
 # Electrons = # Protons
 Mass number – total number of protons and neutrons that make up
the nucleus of an isotope




Found on the periodic table
Mass # = # Protons + # Neutrons
Comes from the nucleus
To find the number of Neutrons: subtract atomic # from mass #
 # Neutrons = Mass # - Atomic #
 # Neutrons = Mass # - # Protons
Counting Atoms
 Isotopes – atoms of the same element that have different masses,
different number of neutrons { mass number – atomic number =
number of neutrons}
 2 Methods for designating Isotoptes
1. Hyphen Notation - mass number is written with a hyphen after the
name of the element – i.e. Hydrogen – 3 (tritium)
2. Nuclear Symbol – shown the composition of the nucleus – i.e.
235 U (the mass number is the top number, the atomic number is
92
the bottom number)
 Nuclide – term for a specific isotope of an element
 Isotopes of Carbon:
 Carbon-12: has 6 neutrons (12 – 6 = 6)
 Carbon-13: has 7 neutrons (13 – 6 = 7)
Counting Atoms

Isotopes of Carbon
Molar Conversions
 Atomic mass – 1 amu (atomic mass unit), exactly 1/12 the mass of a
carbon-12 atom
 Average atomic mass – weighted average of the atomic masses of the
naturally occurring isotopes of an element, in grams
 Mole (mol) – amount of a substance that contains as many particles as
there are atoms in exactly 12 g of carbon-12
 Molar mass – mass of one mole of pure substance [Avogadro’s number =
6.022 x 1023 ], in grams/mole (used as a conversion factor)
 Formula mass – mass of a pure substance, in grams
Molar Conversions
 How to calculate molar mass:
 Get the mass from the periodic table
 Multiply the mass by the number of atoms
 Add up the total mass
 Example:
 Molar Mass of NO₂
 N 1x14 = 14
 O 2x16 = +32

46 g/mole
Molar Conversions
Dimensional Analysis:
 Given x Unknown = Unknown
 Page 85:
 1) What is the mass in grams of 2.25 moles of Fe? (need to get molar mass
from periodic table.)
Molar Conversions
Dimensional Analysis:
 Given x Unknown = Unknown
 Page 85:
 1) What is the mass in grams of 2.25 moles of Fe? (need to get molar mass
from periodic table.)
 2.25 mole Fe x 56 g Fe = 126 g Fe
1 mole Fe
Molar Conversions
 Continue page 85:
 1) How many moles of calcium, Ca, are in 5 g of calcium? (need to get the
molar mass from the periodic table.)
Molar Conversions
 Continue page 85:
 1) How many moles of calcium, Ca, are in 5 g of calcium? (need to get the
molar mass from the periodic table.)
 5 g Ca x 1 mole Ca = 1.125 mole Ca
40 g Ca
Molar Conversions
 Avogadros Number
 Avogadro’s number = 6.022 x 1023
 6.022 x 1023 molecules or atoms or particles = 2 mole
 Page 86:
 1) How many moles of lead, Pb, are in 1.50 x 10¹² atoms of lead? (need to
use Avogadros number to convert)
Molar Conversions
 Continue page 86:
 1) How many moles of lead, Pb, are in 1.50 x 10¹² atoms of lead? (need to
use Avogadros number to convert)
 1.50 x 10¹² atoms Pb x
1 mole Pb
=
6.02 x 10²³ atoms Pb
2.49 x 10ˉ¹² moles Pb