The Periodic Table

• Elements are arranged in order of proton numbers
• Vertical columns of elements are called groups
• Horizontal rows of elements are called periods
• Across the period, the elements change from metals to non-metals
• Elements close to the staircase line, like silicon or germanium, can have some properties of the metals and of the non-metals --> metalloids
• Group number: indicates number of electrons in the outermost shell of an atom
• Period number: indicates the number of electron shells in an atom
• Metals always from positive ions while non-metals always form negative ions
  

Periodic Law

• Elements in same group has the same number of valence shell electrons which the amount is the same as the group number.
• e.g. Group II has elements with valency of 2 electrons.
• The charges relates to the group number and number of valence electrons.
• Elements on left side periodic table lose ions to form cation.
• Elements on right side periodic table gain ions to form anion
• Elements in Group IV can lose or gain electrons depending on reacting element.
• Transition metals may form variable cation of 2+ or 3+
• Elements in same group form same type and number of bonds due to the same number of valence electrons.
• e.g. Sodium in Group I forms NaCl, so other elements in Group I does the same. (RbCl, KCl, LiCl, CsCl)
• From left to right, elements gradually change from metal to non-metal
• Elements close to dividing line in periodic table in back part of the note (in bold) are called metalloids having properties of metals and non-metals.
  

• Ionization energy is the minimum amount of energy required to completely remove an electron from a gaseous atom
• The Electron affinity of a molecule or atom is the energy change when an electron is added to the neutral atom to form a negative ion. (in simple terms, it refers to how much an element loves an electron)
• Electronegativity is the tendency for an atom to attract electrons to itself
• Electropositivity is the tendency for an atom to lose an electron

Changes in Group

• Proton number increase going down the group
• On each sides of periodic table, the change of the proton number small & gradual
• In transition metals, the gradual change is larger
  

Using the Periodic Table

• Given chlorine, iodine and bromine of Group VII forms molecules of Cl₂, I₂ and Br₂ respectively, predict the molecular formula of Fluorine. eg. F₂
• From example, we know elements in same group form same formula.
  

• Properties of element changes down the group.
• i.e. given list of Group 7 elements, predict the properties of astatine.

Octet rule

Cations

Anions

Isotopes

Differences between metals and non-metals

Group I metals

• called Alkali metals - the metals react with water to form alkaline solutions: the solutions turn red litmus paper blue
• have one outer shell electrons
  
• shiny, silvery solids
• soft, easily cut with scalpel
• low densities and melting points - these increases down the group
  
• good conductors of heat and electricity
• reacts easily in air so they're kept in oil
• chemically reactive - reactivity increases down the group (Caesium most reactive of all metals
• This is because the atoms become larger as we go down the group.
  
• It becomes easier for the outermost, single valence electron to escape to form an ion as the attractive force of the nucleus is further away and weaker
• have 1 valence electron
• loses one outermost electron to form an ion of +1 charge eg. Li⁺, Na⁺, K⁺
• react violently with air or oxygen, catching fire, and burning with characteristic flame colors to form white oxides. To avoid this, the metals are stored under oil
• eg. 4K (s) + O₂ (g) ----> 2K₂O (s)
• They also react vigorously with water, forming the alkaline hydroxide and releasing hydrogen gas
• eg. 2Na (s) + 2H₂O (l) ---> 2NaOH (aq) + H₂ (g)
• Because alkali metals are such reactive metals, they combine directly with reactive non-metals such as the halogens to form salts
• eg. 2Na (s) + Cl₂ (g) ----> 2NaCl (s)
  

Flame colour

Reactions of the Alkali metals

Group VII (Halogens)

• reactive non-metals
• have seven outer shell electrons
• poisonous
• low melting and boiling points - increases down the group
• elements become darker and solidify down the group
  
• reactivity decreases down the group
• because of their atomic size, which increases down the group
• therefore it becomes more difficult for the nucleus to attract an electron to form an ion
• most reactive is fluorine; least reactive is iodine
• all halogens form ions with single negative charge eg F^-, Cl^-, Br^-
• Exists as diatomic molecules eg F₂, Cl₂, Br₂, I₂
• reacts vigorously with metals to form ionic salts
  
• halogens become less reactive down the group
  

• Any halogen above another in the group will displace it from a solution of its salt ----> displacement reactions
• This means that the more reactive halogen can take the place of the less reactive halogen of its salt.
• eg. Cl₂ + 2KBr -----> 2KCl + Br₂
• When chlorine gas is bubbled through a colorless solution of potassium bromide, reddish brown color of bromine is seen.
• eg. Cl₂ + 2KI ----> 2KCl + I₂
• When chlorine gas is bubbled through a colorless solution of potassium iodide, it turns brown and finally a black precipitate of iodine is formed.
  

Group O (Noble/Inert Gases)

• Exist as monatomic (single atom)
• least reactive elements in the gaseous state; do not form bonds
  
• low melting and boiling points
• have stable electronic configuration with full electrons on their shells
  
• chemically inert
• because the outermost shell of the element is full.
  
• It does not tend to combine with other elements, either covalently or ionically, and there is chemically inactive
  

Transition metals

• often form colored compounds
• can have variable oxidation states ---> no fixed number of valence electrons
• form complex ions eg MnO₄^- (manganate(VII) ions)
  
• high melting and boiling points
• high densities
• can have catalytic properties
  

Uses of transition metals

• Since transition elements speed up chemical processes in industries, they saves time in manufacture
• Less energy is needed for manufacture in industries, hence lower cost
• Since less energy is needed, more energy resources can be conserved, e.g. oil to generate electricity in producing iron.
  

MCQ