Unit 1: Elements & Compounds

Alkali Metals

soft-siliver-coloured elements: soild a room temp. exhibit metalic properties; react violently in water to liberate hydrogen gas; reacts with halogens to form compound (ie. NaCl); stored under oil or in a vaccum to prevent reaction with air

Alkaline Earth Metals

Light, very reactive; soild at room temp.: exhibit metalic properties; reacts with oxyen to form oxides (ie. MO); all except Be; will react with hydorgen to form hydrides; reacts with water to relase hydrogen

Transition Metals

exhibit a range of chemical and physical properties: strong, hard metals with high melting points; good conductors of electricity; form mutlivalent ions (many react with oxygen to do so); many react with oxygen to form oxides some react with acids to releases hydrogen gas

Lanthanide

elements with atomic numbers 57-70

Actinides

elements with atomic numbers 89-102

Non-metals

located right side of the “step” gain from metals to crete ionic compound: soild, liquid or gas at room temp. 4 or more valence electrons, share with metals to acheive full octet. Covalent or molecular compouds. Can form ionic compunds by gaining elections forming anions (negative ions)

Halogens

may be soilds, liquids or gases at room temp.; (non-metalix properties) not lustrous, and insluatirsm extremtly reactive (espically F); readily with hydorgen and metal

Naturally Occuring

Elements 1-92 (Hydrogen- Uranium)

Synthetic

93- 118 (Neptunium- Oganesson

Radioactive

84 and above and seeral naturally occuing isotopes of lighter elements

Radioactive properties

unstable nuclei and emit hig energentic particles; number of protons chages and atom changes into differnt eements

Transuranium

92 and greater

Alpha (α)

^4 2 He: same structure of **hellium nuclei (2+ charge)** can be **blocked by paper**

Beta (β)

^0 -1e:: negativitly charge **electrons** and can be blocked by ==**aluminum foil**==

Gamma (^0 0γ)

no mass, no charge, and travel at the speed of light, can be blocked by lead, **DNA INTERCOLATION break H Bonds**

Nuclear force

holding nucleaus togther (in atom)

Electromagnectic force

negatively charged elctrons to the postively charged nucleus however, as more elctron shells are added to the atom, attraction weakens

Metals

4 or less valacne electrons, loses them, cations (positive ions) , unable to share compounds with each other

Moving across periodic table

atomc radius decasres and ionization energy increases (becomes progressicily more diffcult to remove electron in gasous state)

Atomic Radius

estimate of the radius of an atom form the centre of the nucleus to the valence shell. Measued by taking the average distance between the nuclei to two identical neighbouring atoms. Measured using

picometers (10^-12)

Atomic Radius Trend

* AR **increases** as you move down a __family/grou__p: more ==protons== arre added to an ==atom== which increases the size of the ==atom== along with more ==electrons shells==. The i==nner elctrons== exhibit a ==screening effect== from the ==protons’ positive pull== onto the ==outer elctons== which allow the outer elctrons to space ==fruther away== from the ==nucleus’ centre.==
* **decreases** as you are moving across a __row/period__: more electrons are added to the ==same electron shel==l, and ==more proton==s are added to the ==nucleus==, increasing the ==nuclear charge==. This will cause the ==valance shell== to have a ==stronger attractio==n to the ==nuclues== pulling them ==tighter==
* Ionic radius of a __cation__ is **smaller** than atomic radius: ==electromagnetic force== pulling on each electron is now shared among ==fewer electron==s,
* Ionic radius of an __anion__ is **larger** that atomic radius: more ==replusion== among electrons

Ionization Energy

energy needed to remove an electron from a gaseous atom (forming a cation). The first electron removed is the electron in the outer she;l of the atom and is the least attracted to the atom. It is possible to remove every electron form the atom however after each sucessive electron is removed the amount of energy required increases. Noble gases have large amounts

Ionization Energy Trend

* IE decreases moving down a family/group: shells being added casuing the screening effect reducing the attraction of the valence electrons to the nucleus
* increases as you move across row/period: nuclear charge inncreases and attraction strenght increases

Electron Affinity

is amount of electron relasesd when an electron is added to a gasesous atom (measuing heat\`\`\`\`\`\`\`\`\`\`\`\`\`\~kb/mol.) A large value indidcates that the trend is likely to occur; sometimes zero as when an electron is added to an atom, its attraction to its nucleus cancels out by replusion of the atoms. Noble gases are 0

Electron Affinity trend

* EA decreases as you move down a group: more shells being added and screening effect redcuing the attraction of an electron
* increases as you move across a period: nuclear charge increases and the attraction increases

Electronegativity

relative strenght of attraction by an atom to from a bonding electron pair in a compound. This is a therotical scale devised by Pauling with highest value is Flourine (4.0). By comparing the EN values of 2 different atoms, can predict if the bond is covalent or ionic between them

Electonegatvity Trend

* EN decreases as you move down a family/group: atom size increases (as more protons are added) More electron shells are added bt the inner sheel creates a sheilding effect (causes the EN to decrease)
* increases as you move across a period/row: nuclear charge increases and the shared electrons in valance sheel are attracted more strongly

Polaraity

affects their intermolecular *attractions*, and solubility. Polar compounds disolve with other polar compounds or ionic; and non-polar compounds disolves no polar. Polar compounds repelled from or attract charged objects. EN affacts shape and size

VSPER THEORY

The approx shpae of a strcutre can be predicited using the valance hseel and its elctrons and lone pairs.