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Standard 1Atomic Structure
Chapters 4-6
Nobel gaseshalogensSemi-metalsTransition metals
Alkaline earth metalsAlkali metals
Metal/non-metalboundary.
Metals
Non-metalsPeriodic Table.
Summary 1
• Which elements are semi-metals?
• Metals:– Good conductors– Solid (except mercury)– Lose electrons– Example = aluminum
• Semi-metals (metalloids):– Have properties of both
metals and non-metals– Common use =
– semi-conductors
– Example = silicon
1b: groups of the Periodic Table
• Non-metals:– poor conductors– Mostly liquid/gas– gain electrons– Example =
nitrogen
• Halogens:– Extremely reactive– Gain 1 electron– Mostly gases– Example = fluorine
Summary 2
1. Describe the differences between metals and non-metals.
2. Give an example of a metal3. Give an example of a non-metal
1c: Periodic Groups• Alkali metals
– Extremely reactive
– Lose 1 electron– Example: sodium
• Alkaline earth metals– Reactive– Lose 2 electrons– Example: calcium
• Transition metals– Can lose different
numbers of electrons
– Example: copper
• Noble gases– Extremely un-
reactive– Gases!– Example: helium
Summary 3
• Which group of metals are most reactive?
• 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.
1a: organization of the periodic table
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
1
3 4 10
2
97 865
6
C12.011
Atomic number
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)
Isotopesex:
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
Na Na+
symbol symbol
Summary 7
• If a magnesium atom loses two electrons, how many electrons will this magnesium ion have?
1 valence e- 4 valence e-
• Valence electrons are:• responsible for chemical behavior of atom • used for chemical bonding• located in the outer orbital
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?
• 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.
1c: Periodic Trends
• General trend for electronegativity:
Increasing electronegativity
Incre
asin
g
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 ionization energy
Incre
asin
g
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)
Decreasing atomic size
Incre
asin
gdecreasing
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?
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
Standard 1e: The structure of an atom
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?