Atomic Structure and The Periodic Table
The Structure of Atoms:
Protons: Positively charged particles
Neutrons: No charge and similar mass to a proton
Electrons: Negatively charged particles, that have very low mass. They are found orbiting the nucleus.
Nucleus: Dense centre of an atom, made up of nucleons (protons and neutrons)
An atom is held together by electrostatic forces of attraction between the nucleus and the electrons.
Element: A type of atom, classified by the number of protons.
Molecule: A substance that consists of two or more atoms that are chemically combined. They can be the same, such as N2 or different such as H2O.
Ion: An atom that has lost or gained electrons in order to reach a stable state, creating a charged atom.
Atomic Representation:
Element name: The name given to a type of atom, such as chlorine or boron.
Atomic Symbol: a letter, written as a capital, or two letters that represent an element, such as Cl or B.
Atomic Number: Z: total number of protons, always a whole number. The periodic table is arranged by atomic number.
Mass Number: A: The number of protons plus the number of neutrons. Always a whole number.
Atomic Mass: Average mass of the isotopes of that element. Usually not a whole number.
Isotopes:
Different forms of the same element. They contain the same number of protons but a different number of neutrons.
They will have very similar chemical properties, such as heat of combustion, reactivity with water and pH but different physical properties such as nuclear stability and mass.
The atomic number stays the same between isotopes of the same element, but the mass number will change.
Carbon has three naturally occurring isotopes: C12 C13 and C14.
Relative Atomic Mass: depends on the relative abundance of isotopes f an element, and each element has only one RAM. It is the mean mass of the naturally occurring isotopes of an element on the RAM scale compared with carbon-12
Dmitri Mendeleev (1834-1907) notice trends in different elements and the chemical properties they presented, and developed an early form of the periodic table. He arranged them by increasing relative atomic mass (not mass number).
Organisation of the periodic table:
Trends:
Ionisation energy: The amount of energy required to remove an electron from an atom in gaseous form. Written in kJ mol-1.
Electronegativity: An element's ability to attract electrons to itself. expressed in arbitrary units on the Pauling scale. Nobel Gases are not given a value as they do not readily form compounds. Moving down a group, electronegativity decreases (valence electrons are further from the nucleus) and increases across the period (nuclear charge increases so greater pull on valence electrons)
Metallic Characteristics: Determined by an elements ability to lose electrons. Decrease as you move across (from metals to non-metals) and increases down (increased number of shells, so ionisation energy increases).
Oxidising and Reducing Strength: Oxidising strength is how easily an element gains electrons. Trends are same as electronegativity. Strong oxidants gain electrons easily and are themselves reduced. Reducing strength trends match ionisation energy.
Reactivity: Is different for metals and non-metals. Metals: increases down and decreases across. Non-metals: Decreases down and increases across (except Nobel gases).
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| Diagram showing parts of an atom from here |
Neutrons: No charge and similar mass to a proton
Electrons: Negatively charged particles, that have very low mass. They are found orbiting the nucleus.
Nucleus: Dense centre of an atom, made up of nucleons (protons and neutrons)
An atom is held together by electrostatic forces of attraction between the nucleus and the electrons.
Element: A type of atom, classified by the number of protons.
Molecule: A substance that consists of two or more atoms that are chemically combined. They can be the same, such as N2 or different such as H2O.
Ion: An atom that has lost or gained electrons in order to reach a stable state, creating a charged atom.
Atomic Representation:
Element name: The name given to a type of atom, such as chlorine or boron.
Atomic Symbol: a letter, written as a capital, or two letters that represent an element, such as Cl or B.
Atomic Number: Z: total number of protons, always a whole number. The periodic table is arranged by atomic number.
Mass Number: A: The number of protons plus the number of neutrons. Always a whole number.
Atomic Mass: Average mass of the isotopes of that element. Usually not a whole number.
Isotopes:
Different forms of the same element. They contain the same number of protons but a different number of neutrons.
They will have very similar chemical properties, such as heat of combustion, reactivity with water and pH but different physical properties such as nuclear stability and mass.
The atomic number stays the same between isotopes of the same element, but the mass number will change.
Carbon has three naturally occurring isotopes: C12 C13 and C14.
Relative Atomic Mass: depends on the relative abundance of isotopes f an element, and each element has only one RAM. It is the mean mass of the naturally occurring isotopes of an element on the RAM scale compared with carbon-12
(% of isotope 1 × mass of isotope 1) + (% of isotope 2 × mass of isotope 2) ÷ 100
Relative Isotopic Mass
Each isotope of each element has its own RIM. It is determined in comparison to carbon-12 (which is exactly 12)
The Periodic Table:
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| Periodic Table from here |
Dmitri Mendeleev (1834-1907) notice trends in different elements and the chemical properties they presented, and developed an early form of the periodic table. He arranged them by increasing relative atomic mass (not mass number).
Organisation of the periodic table:
Periods go across, each is a new electron shell.
Groups go up and down the columns. For 1-2 and 13-18 this relates to the number of valence (outer) shell electrons.
Trends:
Atomic size: Decreases across (more protons attracting the electrons) and increases down (greater number of shells being filled by electrons, and inner electrons shielding the outer from the pull of the nucleus)
Ionisation energy: The amount of energy required to remove an electron from an atom in gaseous form. Written in kJ mol-1.
Decreases down the period (valence electrons are further form the nucleus) and increases across (nuclear charge is increasing, whereas shielding affect is relatively constant)
Electronegativity: An element's ability to attract electrons to itself. expressed in arbitrary units on the Pauling scale. Nobel Gases are not given a value as they do not readily form compounds. Moving down a group, electronegativity decreases (valence electrons are further from the nucleus) and increases across the period (nuclear charge increases so greater pull on valence electrons)
Metallic Characteristics: Determined by an elements ability to lose electrons. Decrease as you move across (from metals to non-metals) and increases down (increased number of shells, so ionisation energy increases).
Oxidising and Reducing Strength: Oxidising strength is how easily an element gains electrons. Trends are same as electronegativity. Strong oxidants gain electrons easily and are themselves reduced. Reducing strength trends match ionisation energy.
Reactivity: Is different for metals and non-metals. Metals: increases down and decreases across. Non-metals: Decreases down and increases across (except Nobel gases).
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