Download Standard 1: Atomic & Molecular Structure

Standard 1
Atomic Structure
Chapters 4-6
Periodic Table.
Non-metals
Metals
Nobel gases
halogens
Semi-metals
Transition metals
Alkaline earth metals
Alkali metals
Metal/non-metal
boundary.
Summary 1
• Which elements are semi-metals?
1b: groups of the Periodic Table
• Metals:
– Good conductors
– Solid (except mercury)
– Lose electrons
– Example = aluminum
• Non-metals:
– poor conductors
– Mostly liquid/gas
– gain electrons
– Example = nitrogen
• Semi-metals (metalloids): • Halogens:
– Extremely
– Have properties of both
reactive
metals and non-metals
– Gain 1 electron
– Common use =
– Mostly gases
– semi-conductors
– Example = fluorine
– Example = silicon
Summary 2
1. Describe the differences between
metals and non-metals.
2. Give an example of a metal
3. Give an example of a non-metal
1c: Periodic Groups
• Alkali metals
– Extremely
reactive
– Lose 1 electron
– Example: sodium
• Transition metals
– Can lose different
numbers of electrons
– Example: copper
• Alkaline earth metals
– Reactive
– Lose 2 electrons
– Example: calcium
• Noble gases
– Extremely unreactive
– Gases!
– Example: helium
Summary 3
• Which group of metals are most
reactive?
1a: organization of the periodic table
• The Periodic Table: organizes elements
in groups and periods.
•
Groups/families: elements have the
same physical and chemical properties.
•
Rows/periods: elements have the
same number of electron shells.
Summary 4
1. Name another element that would have
similar chemical properties to chlorine.
2. Name an atom that is in the same
period as chlorine.
• The Periodic Table: organizes elements
according to atomic number
• Atomic number = number of protons
Atomic number
6
C
12.011
1
3
2
4
5
6
7
8
9 10
Mass
• Mass number: the number of protons
and neutrons in an atom (units = amu)
• Atomic mass (shown on the periodic
table): the average mass of all isotopes
• Isotope: an atom with the same number
of protons and a different number of
neutrons
• Note: atomic mass generally increases
across the periodic table but not always…
(look at atomic number 27&28, 52&53)
Isotopes
ex:
Summary 5
1. What is the mass number for each
isotope of neon shown in the example?
2. What is the atomic mass for neon?
Standard 1d: electrons
• All atoms have an equal number of
protons and electrons
– Atoms are electrically neutral
• Atoms have no charge
• Symbol: Ne
An equal number of
positive protons and
negative electrons
results in zero charge
Summary 6
• How many electrons are in a magnesium
atom?
• When an atom gains or loses electrons it
becomes an ion
– Ion = charged particle
• number electrons ≠ number protons
symbol
Na
symbol
Na+
Summary 7
• If a magnesium atom loses two
electrons, how many electrons will this
magnesium ion have?
•
•
•
•
Valence electrons are:
responsible for chemical behavior of atom
used for chemical bonding
located in the outer orbital
1 valence e-
4 valence e-
Summary 8
1. How many valence electrons does
nitrogen have?
2. How many total electrons does
nitrogen have?
Identifying Atoms by Emission Spectrum:
•Adding energy ‘excites’ electrons.
•Electrons release energy when they return
to the ‘ground state’ (lowest energy level)
•Released energy = ‘emission spectrum’
•Each atom has a unique emission spectrum
•Scientists use this information in many
ways:
•CSI can identify elements in an unknown
sample
•Astronomers can identify elements in
stars across the universe
Summary 9
What causes an emission spectrum?
1c: Periodic Trends
• Electronegativity: The ability of an atom
to attract an electron
• Example: chlorine is very electronegative
because it wants to ______ an electron.
• Example: sodium is not very
electronegative because it wants to
______ an electron.
• General trend for electronegativity:
Increasing
Increasing electronegativity
Note: for noble gases electronegativity = zero
Summary 10
1. Which is more electronegative: iodine
or chlorine?
2. Which is more electronegative: argon
or chlorine?
• Ionization energy: the energy needed to
remove an electron from an atom
• Example: fluorine has a high ionization
energy because it wants to ______ an
electron.
• Example: potassium has a low ionization
energy because it wants to ______ an
electron.
• General trend for ionization energy:
Increasing
Increasing ionization energy
Note: noble gases have a high ionization energy
Summary 11
1. Which has a higher ionization energy:
iodine or chlorine?
2. Which has a higher ionization energy:
argon or chlorine?
3. Which has a lower ionization energy:
chlorine or magnesium?
• General trend for atomic size (volume)
Increasing
decreasing
Decreasing atomic size
Summary 12
• Which is larger: magnesium or calcium?
• Which is larger: magnesium or chlorine?
General trend for ionic size.
• When atoms lose electrons they get
much smaller
• When atoms gain electrons they get
much larger
Summary 13
Why is Na+ smaller than Na?
Standard 1e: The structure of an atom
1. All the mass of an atom is in
the nucleus (Protons &
neutrons are in the nucleus)
2. In between the nucleus and
the electrons there is only
empty space
Summary 14
Which particles inside the atom have
mass?
Earnest Rutherford
Rutherford demonstrated that the entire
atom is 10,000 times larger than the nucleus
• The rutherford experiment:
• A stream of positive particles (alpha particles)
is aimed at a piece of gold foil.
• Only 1 in 8000 particles is deflected (pass
close to the gold nucleus).
• All other particles travel through ‘empty space’
Summary 15
• How does Rutherford’s experiment
demonstrate that an atom is mostly
empty space?