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• Atoms

• Atomic structure

• The Nucleus

• It’s in the middle of the atom
• It contains protons and neutrons
• It has a positive charge because of the protons.
• Almost the whole mass of the atom is concentrated in the nucleus.

• Atoms

• Atomic structure

• The Electrons

• Move around the nucleus
• They’re negatively charged.
• They’re tiny, but they cover a lot of space..
• The volume their orbits occupy determines how big the atom is.
• They have virtually no mass.
• They occupy shells around the nucleus.
• These shells explain the whole of chemistry.

• Atoms

• Atomic structure

• Summary

• Atoms

• Electron Shell Rules

• Electrons always occupy SHELLS or ENERGY LEVELS.
• The LOWEST energy levels are ALWAYS FILLED FIRST.
• Only a certain number of electrons are allowed in each shell:
• 1st shell: 2
• 2nd shell: 8
• 3rd shell: 8
• Atoms are much HAPPIER when they have FULL electron shells.
• In most atoms the OUTER SHELL is NOT FULL and this makes the atom want to REACT.

• 1st

• 2nd

• So, how do we know how many electrons, protons and neutrons there are?

• Atoms

• Atomic Number and Mass Number

• We just need to know these two simple numbers

• Na

• 23

• 11

• THE MASS NUMBER

• THE ATOMIC NUMBER

• - Number of Protons (and electrons)

• The atomic number tells you how many protons there are.
• This also tells you how many electrons there are.
• To get the number of neutrons – just subtract the atomic number from the mass number.
• The mass number is always the biggest number. It tells you the relative mass of the atom.
• The mass number is always roughly double the atomic number.
• Which means there’s about the same number of protons as neutrons in any nucleus.

• - Total of Protons and Neutrons

• Atoms

• Atomic Number and Mass Number

• Also known as the NUCLEON NUMBER (A)

• Also known as the
• PROTON NUMBER (Z)

• This is known as the NUCLIDE NOTATION

• X

• A

• Z

• Each different type of atom is called a NUCLIDE

• ISOTOPES

• Different forms of the same element

• What’s the difference between these two?

• In the nucleus, this one has 6 protons and 6 neutrons.

• In the nucleus, this one has 6 protons and 7 neutrons.

• Isotopes are atoms with the same number of protons but a different number of neutrons.

• Isotopes have the same atomic number, but different mass numbers

• C

• C

• 12

• 6

• 13

• 6

• ISOTOPES

• Different forms of the same element

• What are the features of isotopes?

• Most elements have different isotopes but there’s usually only one or two stable ones.

• The other isotopes tend to be radioactive, which means that they decay into other elements and give out radiation. This is where all radioactivity comes from – unstable radioactive isotopes undergoing nuclear decay and spitting out high energy particles.

Calculation the average relative atomic mass

Radioactivity

• Radioactivity

Learning objectives

• 10.1.2.3 explain the nature of radioactivity and the use of radioactive isotopes

Evaluation criteria

• knows the nature of radioactive radiation
• able to make equations of reactions of radioactive decay.

RadioActivity – What is it?

• All substance are made of atoms. These have electrons (e) around the outside, and a nucleus in the middle. The nucleus consists of protons (p) and neutrons (n), and is extremely small. (Atoms are almost entirely made of empty space!)
• In some types of atom, the nucleus is unstable, and will decay into a more stable atom. This radioactive decay is completely spontaneous. The energy that is released from the nucleus of the atom is radiation.

RadioActivity – What is it?

RadioActivity – Alpha Particles

RadioActivity – Beta Particles

RadioActivity – GAMMA WAVES

RadioActivity – SUMMARY

ISOtope notation

• Isotope Notation includes additional information about an isotope. In addition to the chemical symbol, the mass number and the atomic number are included. This allows information about the nucleus to be determined.

NUCLEAR DECAY REACTIONS

• SOME Elements are naturally unstable
• they tend to undergo alpha-decay or beta-decay to become more stable.
• They may take several steps in order to achieve this, thus we observe decay chains (also known as series decay) for most radioactive elements.
• For example, it takes U-238 14 steps of alpha and beta decay to become completely stable

Identifying alpha and beta decay REACTIONS

Identifying alpha and beta decay REACTIONS

Periodic Table & Trends

History of the Periodic Table

• 1871 – Mendeleev arranged the elements according to: 1. Increasing atomic mass 2. Elements w/ similar properties were put in the same row
• 1913 – Moseley arranged the elements according to: 1. Increasing atomic number 2. Elements w/ similar properties were put in the same column

Group Names

Alkali +1	Alkaline Earth Metals +2		+3		-3	-2	Halogen -1	Noble Gases 0
H 1								He 2
Li 3	Be 4		B 5	C 6	N 7	O 8	F 9	Ne 10
Na 11	Mg 12		Al 13	Si 14	P 15	S 16	Cl 17	Ar 18

S & P block – Representative Elements

• S & P block – Representative Elements
• Metalloids (Semimetals, Semiconductors) – B,Si, Ge, As, Sb, Te (properties of both metals & nonmetals)
• Columns – groups or families Rows - periods

• METALS

• TRANSITION METALS

• NONMETALS

Periodic Groups

• Elements in the same column have similar chemical and physical properties
• These similarities are observed because elements in a column have similar e- configurations (same amount of electrons in outermost shell)

Periodic Trends

• Periodic Trends – patterns (don’t always hold true) can be seen with our current arrangement of the elements (Moseley)
• Trends we’ll be looking at:
• Atomic Radius
• Ionization Energy
• 3. Electronegativity

Atomic Radius

• Atomic Radius – size of an atom
• (distance from nucleus to outermost e-)

Atomic Radius Trend

• Group Trend – As you go down a column, atomic radius increases
• As you go down, e- are filled into orbitals that are farther away from the nucleus (attraction not as strong)
• Periodic Trend – As you go across a period (L to R), atomic radius decreases
• As you go L to R, e- are put into the same orbital, but more p+ and e- total (more attraction = smaller size)

Ionic Radius

• Ionic Radius –
• size of an atom when it is an ion

Ionic Radius Trend

• Metals – lose e-, which means more p+ than e- (more attraction) SO…
• Cation Radius < Neutral Atomic Radius
• Nonmetals – gain e-, which means more e- than p+ (not as much attraction) SO…
• Anion Radius > Neutral Atomic Radius

Ionic Radius Trend

• Group Trend – As you go down a column, ionic radius increases
• Periodic Trend – As you go across a period (L to R), cation radius decreases,
• anion radius decreases, too.
• As you go L to R, cations have more attraction (smaller size because more p+ than e-). The anions have a larger size than the cations, but also decrease L to R because of less attraction (more e- than p+)

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