Topic 2 - Patterns In The Periodic Table PDF

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Topic 2: Patterns in The Periodic Table

Metals      

Solid at room temp. (except Hg) Reflective surface – shiny Conduct heat and electricity Malleable and ductile Lose electrons and form cations (positively charged) in reactions Lower left on the periodic table

Non-Metals     

Found in all 3 states Poor conductors of heat and electricity Solids are brittle Gain electrons in reactions to become anions (negatively charged) Upper right on the periodic table (plus H)

Metalloids    

Some properties of metals and some of non-metals Also known as semiconductors Elements alone diagonal in P group E.g. silicon o Shiny (metal property) o Conducts electricity (metal property) o Does not conduct heat well (non-metal property) o Brittle (non-metal property)

Periodic Table   

Vertical columns: elements with similar chemical and physical properties. Columns called groups/families Horizontal rows (called periods or shells): shows the pattern of properties repeated in the next period.

Neutral atoms o o o o

Protons have a positive charge Electrons have a negative charge Neutrons have no charger Neutral atoms have equal number of electrons and protons

Orbitals     

Electrons are found outside the nucleus in electronic orbitals (electron charge clouds). Orbitals are spaces around the nucleus where the probability of finding electrons is high. Orbitals can have between 0-2 electrons in each. Orbitals consist of s, p, d, and f. All atoms have the same pattern of orbitals whether or not the orbitals are occupied by electrons.

Energy Levels/Shells    

Electronic configuration determines the energy level of each orbital. It is the order in which electrons fill the orbit. Electron shell – region of space where electrons spin around the nucleus Shells further from the nucleus can hold more electrons than those near the nucleus. Sub-levels/sub-shells within the main shells which are called s, p, d, or f orbitals.

Level 1 Shell. - S orbital    

1s Orbital is spherical (max of 2 electrons in it). Lowest energy in atoms. Only orbital in level 1 of all atoms. E.g. Hydrogen electronic configuration: 1s1 Helium electronic configuration: 1s2

Level 2 Shell – S & P orbital     

2s and 2p orbitals are shaped like 2 elongated balloons attached. There are 3 lots of 2p orbitals for every atom (2px, 2py, 2pz). P orbitals hold up to 6 electrons. All orbitals have the same energy. E.g. Carbon electronic configuration: 1s2, 2s2, 2p6 Oxygen electronic configuration: 1s2, 2s2, 2p4

S orbitals are filled before P orbitals

Always do up arrows first

Level 3 Shell    

Has a s level, p level and d level. There are 5 lots of d orbitals. D level holds up to 10 electrons. E.g. calcium electronic configuration: 1s2, 2s2, 2p6, 3s2, 3p6, 4s2 (skip 3d10)

Atomic Spectra   

When heated elements emit light Suggests the process emitting light in the atom is quantised (restricted) The electron in the atom may possess only certain energies

Noble Gases   

Elements that have their valence (outermost) shell completely filled. Most stable elements and are all gases. Found on the far-right hand column.

Molecules & Compounds   

Molecules: Groups of atoms that are chemically combined. Elemental molecule: Two or more atoms of the same element combine. Compound molecule: Two or more atoms of different elements combine.

Allotropes  

Allotropes: different forms of the same element. E.g. O2 Oxygen and O3Ozone.

Compounds     

Compounds: Substances that consist of two or more elements that are chemically combined. Compounds have properties that are different to the elements they contain. Atoms combine to form compound in simple ratios: e.g. 1:1, 1:2, 1:3, 2.2 E.g. H2O – hydrogen and oxygen are both gases, the compound water is liquid. Chromatography - A collection of techniques used to separate compounds using a colouring sequence

Types of Compounds: 



Ionic compounds o Where 1 atom or group of atoms loses or gains an electron. o If it loses an electron it has a positive charge. o If it gains an electron it has a negative charge Covalent Compounds o Atoms share an electron o Covalent bonds are harder to break than ionic compounds

Ionic Bonds        

  

Ionic: Attraction between oppositely charged particles Atoms as ions are held together by electrostatic attraction Metal + non-metal Creates network of alternating positively and negatively charged ions (3-dimensional lattice). Formed by the transfer of one or more electrons from one atom to another. Generally metals transfer electrons to non-metals. Always use the simplest ratio of compound. Properties of ionic compounds: high melting/boiling points, hard, many dissolve in water, solutions conduct electricity, ions in solutions have water molecules around them. Dissociation: Anions and cations become separated from each other. Dissociation occurs when ionic compounds are place in water. Ide = gained an electron so it becomes negatively charged e.g. chlorine becomes chloride

Covalent Bonds

  

Covalent: Electrons shared between two atoms Non-polar covalent: Shared evenly Polar covalent: Shared unevenly

Metallic Bonds 

Metallic bond: Metal cations are in a sea of delocalised mobile electrons that are delocalised.

Ions      

Ions are elements (or molecules) that have gained or lost one or more electrons, and as such carry either a positive or negative charge. Positively charged ions (cations) lose electrons Negatively charge ions (anions) gain electrons Stable configurations have 8 electrons in the outer level for smaller p group elements. Opposite charges attract. Like charges repel.

Ionic Compounds     

Have a high melting & boiling point Hard Many dissolves easily in water Solutions conduct electricity Ion in solution are usually hydrated (have water molecules around them).

Dissociation of Ions in water  

Dissociation = When ionic compounds dissolve in water, the anions and cations become separated from each other. This is called. Precipitate = insoluble solid that emerges from a liquid solution

Valence Electrons     

Valence electrons: accessible electrons (found in the outer most shell) Core electrons: the inner filled shells, inaccessible electrons Valence electrons participate in chemical reactions, core electrons do not Valency: The charge that the atom takes on by either gaining or losing electrons in forming a molecule with another atom/s. E.g. Fluorine gains 1 electron, therefore has one extra electron than proton, and therefore has a -1 charge. F-

Columns on the Periodic Table     

Noble gases on RHS (0) Halogens next left (-1) Alkali metals on LHS (+1) Alkaline earth next right (+2) Transition metals in the middle (d section)

Solubility Rules

Naming Ionic Compounds 



Metal + nonmetal = ionic compound (write metal first, then non-metal) 1. Write the name of the metal (or cation) 2. Write the name of the non-metal (or anion) 3. Change the ending to –ide E.g. Cl = chloride, MgBr2 = magnesium bromide, Al203 = aluminum oxide, selenium = selenide

Covalent Compounds        

Covalent compounds are formed by elements sharing electrons Single Bond = 2 electrons are shared between 2 -> H-H Double Bond = 4 electrons are shared between 2 atoms -> Triple Bond = 6 electrons are shared between 2 atoms Elements in the same column as fluorine only form one single bond to other atoms Elements in the same column as oxygen tend to have two bonds, either a double bond to one other atom, or two single bonds to two other atoms Elements under nitrogen tend to have three bonds Elements under carbon have four bonds in total

Lewis Dot Structures – single bonds     

Show electrons in valence shell as dots Write single bond as H:H (hydrogen-hydrogen bond) Show as H. (H with a dot) By sharing electrons each atom is able to fill its outside shell Any chemicals that have lone pair electrons become a base?

Stable Octet Rule  

Sharing of electrons between atoms with only S and P orbitals Compound is stable when 8 electrons are provided around the central atom

Double Bonds 

O2 is a double bond (each O shares 2 electrons, thus a total of 4 electrons are shared)

Triple Bond 

N2 (each N atom shares 3 electrons) = Triple Bond

Ordering Elements in a Molecular Formula  



Metals are written first, e.g. MgO Non-mentals are written in the order below o C, P, N, H, S, I, Br, Cl, O, F o E.g. CH4, Nh3, So2, HF, PCI3 Exceptions o H20, but not NaOH (Na interacts with the O not the H) o HCN not CNH (H-C=N not C=N-H)

Polarity   

 

Non-polar = no separation of positive and negative charge Polar = Charges are separated but electrons are still shared Charge separation = bonds are polar but the molecule is not due to its 3-dimensional shape. It can also occur when there are atoms with different electrons attracting abilities (electronegativity) Water is a bent molecule = this is because it has 2 lone pairs of electrons housed in the orbitals Water is a polar covalent molecule = O atom is very electronegative so takes on a partial negative charge, whilst each H atom takes on a partial positive charge.

Dots are lone pairs of electrons

Water as a solvent

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    

One side of water is negatively charged because the oxygen atom keeps the shared electrons longer than the hydrogen atoms. As a result the oxygen side is negatively charged and the hydrogen side if water is positively charged. NACL: since unlike charges attract, the negative end of water will be attracted to a positive NA. the positive end of water will be attracted to a negative Cl. Interactions between molecules are different between polar and non-polar compounds These interactions affect the physical properties of the compound e.g. melting/boiling point Polar Molecules = one type of interaction is Hydrogen Bonding Polar molecules have an electrostatic interaction – negative part of the molecule is directed towards the positive part of another molecule

Ionic / Covalent Bond (polar/non-polar) Continuum

Polyatomic Ions    

Ions are single atoms that have either gained or lost electrons Ions are molecules that have either positive or negative charge Bonding within ions is covalent – so when it dissolves in water it does not dissociate into separate atoms; the entity has a net charge Examples include ammonium (NH4+), hydroxide (OH-), nitrate (NO3-), sulfate (SO42-), phosphate (PO43-) and carbonate (CO32-).

Polyatomic Ions 



While the bonding within the polyatomic ion is covalent, the charge that is carried by the ion must be cancelled out. This is achieved by forming an ionic bond to an ion of opposite charge Examples:

Re

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

member: o Ammonium -> NH4+ o Carbonate -> CO32+ o Hydrogen carbonate -> HCO3o Hydroxide –> OHo Nitrate -> NO3o Phosphate ->PO43o Sulfate -> SO42o Sulfite -> SO32When any polyatomic ionic dissolves, the polyatomic ions remain intact, but the cations and anions separate.

Mixtures   

The molecules in mixtures are not chemically bonded together -> they maintain their own unique properties The properties of the combined mixture MAY be different to the bulk phase Mixtures can be separate into three types o Solutions (homogeneous mixtures) o Colloids o Suspensions (heterogeneous mixtures)

Solutions      

Homogeneous – uniform throughout. They can be gases, liquids or solids E.g. Air (it contains N2, O2 and a range of other gases) it is the same throughout the world Solvent = The substance that is most present (most abundant) in a solution Solutes = anything that is less present in the solution Solutions will not unmix

Alloys      

A mixture of metals that has different (improved) properties from the metal elements that make it. Homogeneous mixture existing in one solid phase. It is a solid solution Made by melting the ingredients and mixing them together then allowing the mixture to cool and solidify E.g. Steel – an alloy of Fe and C - the presence of the C atoms lowers the ductility where Fe increases in hardness and tensile strength E.g. stainless steel – alloy of iron, chromium, nickel and carbon - strong and hard, resistant to corrosion (due to chromium)

Suspensions   

Heterogenous mixtures – sometimes with visible solutes E.g. coke – bubbles in the coke make it heterogenous It will separate/un-mix

Colloids      

Colloids = emulsions The solute particles must be larger than the size of a molecule but smaller than what can be seen by the naked eye. Colloid particles are so small that they do not settle out Colloids are heterogeneous mixtures that often appear milky or translucent Colloids scatter light Colloids an change reversibly from a fluid to a more rigid (gel) state

Differences between mixtures and compounds Mixtures:    

No chemical bonding occurs between the components of a mixture The properties of the atoms and molecules are unaffected by other components of the mixtures Components of a mixture can be separated by physical means (filtering/evaporating) Mixtures can be both homogenous and heterogeneous

Compounds:  

Can be separated into their constituents only through a chemical reaction Compounds are homogeneous

Physical methods of separation for mixtures          

Evaporation – boiling of a solvent Filtration – separating a suspended solid from a solvent by pouring through a sieve (semipermeable membrane) Fractional distillation – collecting different liquids on the basis of their boiling points Fractional crystallisation – separation of solids based on different solubilities Magnetism – separation of magnetic and non-magnetic species Centrifugation – separation of different species based on density, by spinning a mixture very quickly Winnowing - separation based on weight Adsorption - partitioning two species by selectively adsorbing one onto a surface. Decantation – pouring the less dense liquid from above a denser liquid in a heterogeneous mixture. Floatation - use of froth that separates hydrophobic and hydrophilic species. Used in mineral processing....