Chapter 4
Atoms & The Periodic Table
Physical Science
Jones
What will we be learning?
 Section 1 – Atomic Structure
 Explain Dalton’s Atomic Theory, and describe why it was more
successful than Democritus’ theory.
 State the charge, mass, and location of each part of an atom
according to the modern model of an atom.
 Compare and contrast Bohr’s model with the modern model of
the atom.
Inquiry learning!
 Read pages 104-107.
 Write down important names, dates, and discoveries of parts
of the atom.
 On a piece of copy paper, draw a timeline and include the
important names, dates, and experiments in order. Be sure
to include information about the following people and their
discoveries:
 John Dalton
 Democritus
 Niels Bohr
Atomic Theory
 What did the scientists think the atom looked like?
 It was too small to see
 There was not a lot of technology to “look” at atoms
(microscopes, etc.)
400 BC: Atomic theory
 The idea of matter existing as particles
was introduced as early as 400 B.C.
 Democritus – atoms are “indivisible”
 Aristotle – all matter was continuous
 Aristotle’s proposal lasted about 2000
years
1808: Dalton’s atomic theory
 John Dalton, 1808 – proposed an explanation for the
previous laws:
 1. All matter is composed of extremely small particles called
atoms
 2. Atoms of a given element are identical in size, mass, and
other properties
 3. Atoms cannot be subdivided, created, or destroyed
 4. Atoms of different elements combine in simple, wholenumber ratios to form compounds
 5. In chemical reactions, atoms are combined, separated, or
rearranged
~100 years after Dalton
 Scientists decided that Dalton’s #3 was incorrect: “Atoms
cannot be subdivided.”
 We now know there are subatomic particles (parts of an
atom)
 Proton
 Neutron
 Electron
1913: Niels Bohr
 Danish Scientist
 Said electrons move around the nucleus (where the protons
and neutrons are) like planets orbiting the sun
1925
 Scientists believed that electrons moved more like waves than
particles that orbit the nucleus
Modern Atomic Theory
 What do we believe about what the atom looks like now?
 Dalton’s atomic theory + subatomic particles
 Atoms have a nucleus
 Cluster of protons + neutrons
 Atoms have an electron cloud
 Electrons surround the nucleus in various energy
 Are in motion
 Location of an electron cannot be determined
levels
Energy levels
 Electrons are found in energy levels around the nucleus
 Often described like rungs of a ladder
 Outermost energy level has a certain number of electrons
called valence electrons
 Atoms can lose or gain valence electrons to become more
stable
 Atoms that have lost or gained electrons are called ions
Good Afternoon!
 Please review last week’s content by completing the
following:
 P 110 #1-6
 P 119 #1-5
What will we be learning?
 Section 2 & 3 – The Periodic Table
 Relate the organization of the periodic table to the arrangement
of electrons within an atom.
 Explain why some atoms gain or lose electrons to form ions.
 Determine how many protons, neutrons, and electrons an atom
has, given its symbol, atomic number, and mass number.
 Describe how the abundance of isotopes affects an element’s
average atomic mass.
The Periodic Table
 The elements are not put on the table in a random order.
 Listed by atomic number
 Arranged by number of valence electrons, reactivity, and other
characteristics
 Period = row
 Group = column
 Elements in the same group have the same characteristics
because they have the same number of valence electrons
# of Protons: ________
# of Electrons: _______
____________________
Atomic Number
# of Neutrons: _______
_______________________
Element Name
_______________________
Mass Number
The Periodic Table
 # of protons = atomic number
 # of electrons = # of protons (because the atom is neutral)
 # of neutrons = rounded mass number - # of protons
 Classwork!
 Complete the worksheet – use the periodic table to indicate the
symbol and number of subatomic particles for each element.
Ions
 Ions – atoms of elements that have either lost or gained electrons
 Cation – a positive ion that has lost electrons
 Anion – a negative ion that has gained electrons
 Atoms lose or gain electrons to become more stable. Each atom gains
enough to have 8 valence electrons, or loses enough to have a full
outer energy level.
 Ca: 2 valence electrons – lose them, or gain 6?
 Lose them = Ca+2
 N: 5 valence electrons – lose them, or gain 3?
 Gain 3 = N-3
Isotopes
 Isotopes - Atoms of an element with a different number of
neutrons
 Different number of protons = not ok! (different element)
 Different number of electrons = not ok! (ions)
 Elements exist in different forms with different masses on
Earth.
 Example:
 H-1: “Hydrogen-1” with 1 proton, 1 electron, 0 neutrons.
 H-2: “Hydrogen-2” with 1 proton, 1 electron, 1 neutron.
 H-3: “Hydrogen-3” with 1 proton, 1 electron, 2 neutrons.
Isotopes
 The mass of each isotope is calculated in to the mass number
found on the periodic table
 The average atomic mass is the average of the mass
numbers and their percentage of abundance on earth.
 The average atomic mass is an average mass of ALL THE
ATOMS of that element found on earth.
 Unit: atomic mass units (amu)
The Periodic Table
 Three groups of elements:
 Metals
 Nonmetals
 Semiconductors (also called Metalloids)
Parts of the Periodic Table
 Label and color each of the following:
 Alkali Metals – Group 1 EXCEPT FOR HYDROGEN
 Alkaline-earth Metals – Group 2
 Transition Metals – Groups 3 – 12 (middle of the table)
 Nonmetals – Hydrogen, Groups 17-18, O, S, Se, N, P, and C
 Halogens – Group 17
 Noble Gases – Group 18
 Semiconductors/metalloids – B, Si, Ge, As, Sb, Te
 The rest are just called “other metals”
Parts of the Periodic Table
 Metals
 Conductors of electricity
 Shiny
 Can be made into wires
 Alkali metals are most reactive
 Alkaline-earth metals are found in the human body
 Transition metals are less reactive but very useful
Parts of the Periodic Table
 Nonmetals
 Cannot conduct electricity or be made into wires
 Can form many compounds, especially carbon
 Plentiful on earth
 Halogens are very reactive
 Noble gases do not react at all
Parts of the Periodic Table
 Semiconductors/Metalloids
 Have some properties of metals and some properties of
nonmetals
 Are very useful
 Silicon is most familiar
 Makes up sand (SiO2)
 Found in 28% of the mass of Earth’s crust
 Used in computer parts
What will we be learning?
 Section 4 – Using Moles to Count Atoms
 Explain the relationship between a mole of a substance and
Avogadro’s constant.
 Find the molar mass of an element by using the periodic table.
 Solve problems converting the amount of an element in moles
to its mass in grams, and vice versa.
Counting Units
 What units are used in chemistry to measure a substance?
 Grams (mass) – measured by a scale
 Moles (a large counting unit) – cannot be measured by a scale,
but can be calculated using mass
 A small number with a large unit
 Atoms (a small counting unit) – cannot be measured by a scale,
but can be calculated using moles
 A large number with a small unit
Mass, in grams
 Mass – a measure of the amount of matter
 Molar Mass – the mass of 1 mole of a substance
 The molar mass of any element = its mass number
 How many grams are in 1 mole of Carbon?
 How many grams are in 1 mole of Aluminum?
Mass, in grams
 Mass – a measure of the amount of matter
 Molar Mass – the mass of 1 mole of a substance
 The molar mass of any element = its mass number
 How many grams are in 1 mole of Carbon? 12.0107 g
 How many grams are in 1 mole of Aluminum? 26.981538 g
Mass, in grams
 Mass – a measure of the amount of matter
 Molar Mass – the mass of 1 mole of a substance
 The molar mass of any element = its mass number
 How many grams are in 1 mole of Carbon? 12.0107 g
 How many grams are in 1 mole of Aluminum? 26.981538 g
 The molar mass of any compound = mass numbers added
together
 How many grams are in 1 mole of H2O?
 How many grams are in 1 mole of Na2SO4?
Mass, in grams
 Mass – a measure of the amount of matter
 Molar Mass – the mass of 1 mole of a substance
 The molar mass of any element = its mass number
 How many grams are in 1 mole of Carbon? 12.0107 g
 How many grams are in 1 mole of Aluminum? 26.981538 g
 The molar mass of any compound = mass numbers added together
 How many grams are in 1 mole of H2O?
 1 + 1 + 16 = 18 g
 How many grams are in 1 mole of Na2SO4?
 23 + 23 + 32 + 16+ 16+ 16+ 16 = 142 g
Avogadro’s Constant
 1 mole of ANYTHING = 6.022 x 1023 particles
 = 602,200,000,000,000,000,000,000 particles
 Particles = atoms = molecules
 How many atoms are in 1 mole of Carbon?
 How many molecules are in 1 mole of Na2SO4?
Avogadro’s Constant
 1 mole of ANYTHING = 6.022 x 1023 particles
 = 602,200,000,000,000,000,000,000 particles
 Particles = atoms = molecules
 How many atoms are in 1 mole of Carbon?
 6.022 x 1023 atoms
 How many molecules are in 1 mole of Na2SO4?
 6.022 x 1023 molecules
Conversions
 A calculation that changes one unit to another unit
 Mass  Moles
 Use the molar mass of the substance in the problem
 (How many grams are in 1 mole of _______?)
 P 133 #1-4
 P 134 #1-9
 Turn it into the drawer when you are finished.
Good Afternoon!
 Honors Physical Science Bellwork –
 P 133 #1-4, P 134 #6-9
 Converting between mass and moles
 Hint: given is moles, answer is in grams– multiply given by the molar
mass
 Hint: given is grams, answer is in moles – divide given by the molar mass
 Check your answers with me before you turn it in. Be sure to
show all work and circle your answers!
Chapter 4 Wrap-Up
 Do P 136 #1-18, 20-21, 31
 Turn into the drawer when you are finished.