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PERIODIC TABLE
Chapter 5
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ORGANIZING THE ELEMENTS
Section 1
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LET’S REVIEW!
• Chemical properties
• Any property that can only be tested by changing the
chemical make-up of the substance.
• Physical properties
• Any property that can be tested without changing the
chemical make-up of the substance
• Atomic mass
• Mass of protons and neutrons
• Atomic number
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• Unique to each element, same as number of protons
DMITRI MENDELEEV
• 1870: 63 elements known to
man
• He organized them in order
of their atomic mass, and
saw a pattern from their
properties.
• Was working on this while
Thomson and Rutherford
were still “exploring” the
atom
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DMITRI MENDELEEV
• Arranged his table with repeating properties in
columns, starting a new row each time the
chemical properties repeated
• Left blank spaces in his table, concluding that
these spaces were elements that hadn’t been
discovered yet.
• Based on the patterns and the other elements
around the blank space, he predicted the
properties of those elements
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AN EXAMPLE
What he called ekasilicon – it was discovered a few years later
Prediction
Germanium
Atomic Mass
72 amu
72.6 amu
Density
5.5 g/mL
5.3 g/mL
Appearance
dark gray metal
gray metal
Melting Point
high melting point
937o C
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MENDELEEV’S TABLE
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HENRY MOSELEY
• Mendeleev’s table worked because as
protons increase, atomic mass should
increase, but if there are fewer neutrons
it could decrease
• Errors arose because he was arranging
the table with the wrong number
• 1910: Discovered atomic number and
rearranged the periodic table using this
number, it fell into perfect order
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PERIODIC LAW
• Periodic Law: physical and chemical properties
of the elements are periodic functions of their
atomic numbers
• In other words, when the elements are
arranged by their atomic numbers, you
should see chemical and physical properties
repeating themselves
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ROWS
• Left to right – called periods
• Elements in the same periods
show patterns left to right conductivity/reactivity change,
elements become less metallic
• The period # indicates the
number of energy levels each
atom in the row has
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COLUMNS
• Top to bottom – called groups
• Elements in a group have
similar chemical properties
and show trends top to bottom
• The elements in the same
group (column) have the same
number of valence electrons
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EXPLORING THE PERIODIC
TABLE
Section 2
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REMEMBER
• The periodic table is organized by
atomic number
• For a neutral atom, the number of
protons equals the number of
electrons
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VALENCE ELECTRONS
• The trends found in the
periodic table are a result of
electron arrangement,
specifically, the number of
valence electrons
• Valence Electron: electrons
in the outermost energy
level
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VALENCE ELECTRONS
• The group number of an element will tell you the
number of valence electrons it has
• Group 1: 1 valence electron
• Group 2: 2 valence e- ’s
• Skip 3-12
• Group 13: 3 valence e- ’s
• Groups 14-18: 4, 5, 6, 7, and 8 valence e- ’s respectively.
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ION
• When a neutral atom gains/loses electron(s)
through bonding, the atom is no longer neutral
and has an excess charge
• It becomes an ion
• Ion: a charged atom
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ION
• All atoms want 8 valence electrons in the outer
shell – this makes the shell full and the atom stable
• Elements close to having 8 tend to be the most
reactive.
• Elements already “full” are considered inert,
they don’t react because they don’t need to
• The number of electrons an atom can gain or lose
is equal to the number of valence electrons it has
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LET’S PRACTICE
• Group 16
Give Up? or Gain?
• Group 13
Give Up? or Gain?
• Group 15
Give Up? or Gain?
• Group 2
Give Up? or Gain?
• Group 1
Give Up? or Gain?
• Group 17
Give Up? or Gain?
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ION
• Protons = positive charge
• Electrons = negative charge
• p+ # CANNOT change, but e- #
can
• So…
• If an atom GAINS electrons, is it
more positive, or negative?
• If they LOST electrons?
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ION
• Cation: atoms that LOSE electrons, becoming
more positive
• Anion: atoms that GAIN electrons, becoming more
negative
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ION
• How do we know if an atom is an ion?
• Cations have a +, and anions have a –
superscript
• If an atom has gained 3 electrons
• It has 3 MORE negative particles than positive
particles, it is more negative = Al3-
• If an atom has lost 3 electrons
• It has 3 LESS negative particles than positive
particles, it is more positive = Al3+
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THE PERIODIC TABLE
Section 3
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THE PERIODIC TABLE
• Divided into three major categories based on
general properties
• Metals, Nonmetals, Metalloids (semiconductors)
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METALS
• Like to give up valence
electrons
• Physical Properties: high luster
(shiny), conductive (heat and
electricity), malleable
(bendable), ductile
(stretchable), high density, high
melting point
• Chemical Properties: Most will
react with oxygen
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NONMETALS
• Like to gain electrons
• Physical Properties: dull, don’t
conduct, brittle, low density,
low melting points
• Can be solid, liquid or gas at
room temperature depending
on the element.
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METALLOIDS (SEMICONDUCTORS)
• Share properties of both metals and nonmetals
• Can be shiny or dull, conduct ok, ductile and
malleable or brittle
• These elements have become really important
because of the computer revolution
• Computer chips are made out of semiconductors
(normally Si)
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FAMILIES
• The periodic table can be
further broken down into
families
• Families of elements have
similar properties because
they have the same number
of valence electrons
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FAMILIES
• Metals: Groups 1-12
• Alkali Metals, Alkaline-Earth Metals, Transition Metals
• Groups 13-16 contain both nonmetals/metalloids
• Nonmetals: Oxygen, Nitrogen, Carbon, Sulfur,
Phosphorus, and Selenium
• Metalloids: Boron, Silicon, Germanium, Arsenic
Antimony, and Terellium
• The group is named by the first element in the column
• Nonmetals: Groups 17-18
• Halogens, Noble Gases
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HYDROGEN
• Hydrogen is in group 1 but
is not an alkali metal,
because it is only 1 proton
and 1 electron (no neutrons)
• Its properties are closer to a
nonmetals than to a metal
• it is a colorless, odorless,
explosive gas with oxygen
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GROUP 1: ALKALI METALS
• (excluding H), 1 valence e• Very reactive, especially with water
• Soft, shiny white metals (can be cut with a
knife!)
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GROUP 2: ALKALINE-EARTH METALS
• 2 valence e• Not as reactive as alkali, but
still very reactive.
• Magnesium is used in
flash bulbs
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GROUPS 3-12: TRANSITION METALS
• 1 or 2 valence e• Most are silver and not that
reactive so they have more
everyday uses.
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GROUPS 3-12: TRANSITION METALS
• Two bottom rows, or innertransition
metals
• Lanthanide Series: also called rareearth metals
• Actinide Series: very radioactive
and not easily found in nature
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GROUPS 13-16
• Boron Group: Group 13, 3 valence e• Aluminum is most common and
abundant element on the planet.
• Carbon Group: Group 14, 4 valence e• Pure carbon can be diamonds, soot, or
graphite, silicon and germanium are
used for computer chips
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GROUPS 13-16
• Nitrogen Group: Group 15, 5
valence e• Nitrogen makes up 78% of the air,
Phosphorus is in soaps, and Arsenic
is a well known poison
• Oxygen Group: Group 16, 6 valence
e- ’s
• Oxygen makes up 21% of the air
and is necessary for things to burn
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GROUP 17: HALOGENS
• 7 valence e• All nonmetals (can be solid,
liquid or gas)
• Extremely reactive with alkali
metals
• “Chlorine” added to pools as
a disinfectant is a compound
containing Chlorine, by itself
chlorine is a green gas
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GROUP 18: NOBLE GASES
• 8 valence e- ’s (except Helium)
• Full outer shell of electrons
• All are gases and extremely nonreactive (inert) and found in the
atmosphere
• “Neon” lights contain a variety of
Noble Gases
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