Gr.11 Chemistry - Unit 1 PDF

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Matter, Trends, & Chemical Bonding Periodic Table Development; !

Initial attempts to organize elements used a variety of strategies.

1. Dobereiner - grouped elements of 3 with similar properties together (only in groups of 3). This was called The Law of Triads 2. New-Lands - arranged elements into grouped with similar properties (differed in atomic mass by a multiple of 8. He was the first to assign atomic numbers. 3. Mendeleev - arranged elements by their atomic number and left spaces for future elements to be discovered. He is credited for developing the modern Periodic Table.

Atomic Model Development: 1. Dalton (1803) - Particle Theory: all matter is made of very small invisible particles (atoms which are always in motion). 2. J.J.Thompson (1897) - Raison Bun Model: atoms are composed of smaller parts with charges; localized, negatively charged particles are embedded in a positively charged environment. 3. Rutherford (1911) - Rutherfords Atomic Model: used the results of the Gold Foil experiment. An atom has a very small, dense, positively charged nucleus with electrons in the ‘empty space’ surrounding the nucleus.

nucleus particle: + proton particle: + ∴ they repel each other which gives th deflection. 4. Bohr (1913) - Bohr Rutherford Diagram: expanded on Rutherfords model fffff by proposing that electrons occupy specific energy levels (or orbits) around fffff the nucleus.

Isotopes: • atoms of the same elements but have different atomic masses. • this happens because they have different numbers of neutrons for the element.

Applications of Isotopes: • smoke detectors (Aemericium - 241) • irradiation of food by gamma rays to kill bacteria (Cobalt - 60)

Radio Isotopes: • unstable isotopes (will spontaneously decay, break down into another isotopes or atom) • as isotopes decay they give off particles and energy • half life is a measure of the isotopes stability (half-life = the time it takes for half of the atoms to decay). • Example; a 25.0g sample of 32P is left for 56 days. What mass of this isotope remains? (assuming the half-life is 14 days). ! 0 days →! 14 days ! →! 28 days ! →! 42 days →! 56 days ! 25 g! ! 12.5g !! ! 6.25g! ! ! 3.125g! 1.5625g ∴ it turns into 22/16 S from 32/15 P

Atomic Notation: this is a standard method of displaying information about isotopes.

Average Atomic Number: ← on periodic table Carbon has a mass of 12.011 ← this is an isotope of Carbon ← this is common this is a rare isotopes of Carbon because it has a mass of 13.00→

! ! !

•Isotope abundance = % of atoms that have a specific mass •Example: calculate the average atomic mass of Boron: Atomic Masses:! ! ! ! Abundance: 11 B → 10.0129u!! ! ! 19.9% 12 B → 11.0093u!! ! ! 80.1%

Average Atomic Mass = ∑ (atomic mass X abundance %)+(atomic mass X abundance %)

! !

! !

! !

= (10.0129 X 0.199) + (11.0093 X 0.801) = 10.8110u

Lithium - Li Atomic #: 3 Protons: 3 Electrons: 3 Neutrons: (7-3) = 4 Atomic Mass: 7

Electron Arrangement: • Bohr-Rutherford Diagrams - nucleus - energy levels - electrons

• Lewis Diagrams - symbols - valence electrons (outer most electrons)

• Electron Arrangement (including sub-levels) - a description of location of electrons in an atom Energy Level (n)

# of Sub Levels

Total Electron Capacity

1

1

S-2

2

2

S - 2, P - 6

3

3

S - 2, P - 6, D - 10

4

4

S - 2, P - 6, D - 10, F - 14

Formation of Ions & Electron Arrangement: Ion - a charged atom (number of protons ≠ the number of electrons) • loses or gains electrons to reach a more stable electron arrangement (a stable octet)…a full energy level.

Trends: • Atomic Radius (size of the atom) - amount of space that the electrons take up - within a group the more energy shells = greater the size of the atom - a group is the vertical columns - within a period (this means within an energy level) as more electrons are added, there is also an increase in protons in the nucleus; so there is a greater pull on the electrons (because + and - are attracted to each other), therefore the atoms get smaller as you go right across the periods. and get bigger as you go down the groups • Ionization Energy - the energy required to remove an electron and produce an ion - first ionization energy for the first e- removed, second ionization energy is for the second e- removed. - ionization energy is greater if there is a greater hold on the valence electrons. Electronegativity • - the ability to attract shared e- in a covalent bond - similar factors involved compared to ionization energy • Reactivity - how easily an element will react - metals react by losing electrons - easier that a metal loses electrons; the more likely it will react - low ionization energy = most reactive - bottom left = most reactive - non-metals react by gaining or sharing electrons - stronger hold on electrons = more reactive - top right = most reactive (not including the noble gases)

Ionic Bonds: • are the attraction between a positively charged ion (a cation) and a negatively charged ion (an anion) the ions combine to produce neutral molecules overall (so charges are balanced) • • positive ions form when atoms lose electrons, resulting in full outer shell of electron; usually a metal • negative ions form when atoms gain electrons, also resulting in a full outer shell of electrons, usually a non-metal Example;! ! ! ! ! ! Lewis Structure; ! ! Na → Na⁺ + Le⁻! ! ! ! ! ! [NA]⁺ ! ! Cl + Le⁻ → Cl⁻! ! ! ! ! ! [ Cl ] ⁻

Covalent Bonds: • occur when two atoms share a pair of electrons; with both nuclei pulling on a pair of electrons, the two atoms are held together usually occurs between 2 non-metals (neither atom will give up an electron) • • in a lewis dot diagram, a single unpaired electron is normally a site where a bond will occur • atoms will share electrons in order to achieve a full outer energy level (in the S and P sub levels) this is called the Octet Rule (8 valence e⁻) Example; ! ! H· · H → H - H (a line represents 2 electrons being shared)

Multiple Bonds; • in some cases, sharing one pair of electrons does not result in achieving the octet rule • at times, atoms will share 2-3 pairs of electrons; these are called double bonds or triple bonds.

Polyatomic Ions: • a group of atoms that are covalently bonded together, but have an overall charge these are a little tricker than the other structures • • HINTS: • count valance electrons • then arrange the atoms (first element in the middle) • then arrange electrons to reach a stable structure

Molecular Compounds: • Non-polar covalent bonds - bonding electrons shared equally between two atoms - no charges on atoms Polar covalent bonds • - bonding electrons shared unequally between two atoms - partial charges on atoms • Ionic Bonds - complete transfer of one or more valence electrons - full charges on resulting ions

The Nature of Covalent Bonds: • we know that covalent bonds involve the sharing of a pair of electrons…but what if they don’t share equally? we can determine how equally the electrons are shared by looking at the difference in • electronegativity values for the two atoms involved in the bond the closer the ENEG values are, the more equally they are sharing • • Example; Cl — N — Cl ! ! | ! ! ! ! Cl ! ! ! ! ! ! !

! ! !

ENEG of N = 3.0 ENEG of Cl = 3.2 difference in ENEG = 0.2 ∴ this is a polar covalent bond

Nomen Clature: • is the naming of chemical compounds 1. Binary ionic compounds: • between a metal and a non-metal • metal, nonmetal “ide” 2. Molecular compounds: • use prefixes to indicate the number of atoms present in a molecule • ‘mono’ is not necessary if one atom of the first element is present • non-metal, non-metal “prefix”

• • • • • • • • • •

mono di tri tetra penta hexa heat octa mona deca

3. Multi-valent ionic compounds: • involves a metal and a non-metal, but metal is multivalent name the as a binary ionic compound but most indicate the valence (or charge) of the metal using a roman numeral.

• Example; • Iron (Ⅲ) Oxide → Fe2O3 • Copper (Ⅱ) Chloride → CuCl2 • Iron (Ⅱ) Oxide → FeO

4. Polyatomic compounds • contain at least one ion that is a polyatomic ion (multiple atoms that make up a single ion); recognized as a compound with more than two elements • metal polyatomic ions…usually because most polyatomic ions are anions.

• Example; • Lithium Nitrate → LiNO3 • Magnesium Chlorite → Mg (CIO2)2 • Cobalt (Ⅲ) Sulfate → CO2 (SO4)3

• there are families of polyatomic ions that are similar but have varying numbers of oxygen atoms

• Example; • CIO3⁻ → chlorate (base ion in family) (1 less oxygen) • CIO2⁻ → chlorite • CIO⁻ → hypochlorite (2 less oxygen)

5. Classical System (for multivalent metals) • this system is still used although it doesn’t follow standard naming rules Ending: • use latin names for elements and the ending indicates the higher or ! ic → higher valenc lower valence option ! ous →lower valenc • Iron → ferrum • Tin → stannum • Lead → plumbum • Copper →cuprum • Example; • ferrous chloride → FeCl2 • stannic sulfite → Sn(SO3)2

Acids: 1. Binary acids • contain H plus one other elements • all acids are aqueous (dissolved in water) • aqueous hydrogen element “ide” or hydro ________ -ic acid • • Example; • hydrobromic acid HBr (aq) • hydrochloric acid HCl (aq) 2. Oxyacids • contains H, O, plus one other element • based on H plus a polyatomic ion • aqueous hydrogen polyatomic ion or • ______ -ic acid (if an ‘ate’ ion) • ______ -ous acid (if am ‘ite’ ion) • “ic - ate” “ite - ous” • Example; • chloric acid → HCIO3 • nitrous acid → HNO2 • sulphuric acid → H2SO4

**Only two elements in the formula**

**3 elements in the formula**

Hydrates • ionic compounds that have water incorporated into their crystal structure in a specific ratio • name as regular ionic compound • Example; • Pb(Cl4)2 · 2H2O → Lead (Ⅱ) perchlorate dihydrate • MgCrO · 5H2O → Magnesium cromate pentahydrate • barium hydroxide octahydrate → Ba(OH)2 · 8H2O

Polarity of Molecules: • we can determine whether molecules are polar or not based on two factors: • symmetry of molecules • polarity of bonds in molecules if a molecule has only non-polar bonds, it will be non-polar • • Example; Cs2, H2 • if a molecule has polar bonds and is symmetrical, it will be non-polar • Example; CCl4, CO2 • if a molecule has polar bonds and is not symmetrical, it will be polar • Example; H2O, CHCl3

The Impact of Molecular Polarity: • the polarity of molecules influences the properties of a compound because it has an impact on how molecules interact with each other

Melting Point and Boiling Point: • in the solid state, molecules are close together and have interactions with adjacent molecules • as the sample warms ups, the molecules gain enough energy to loosen their interactions, until the substance melts (turns into a liquid); the stronger the interactions between molecules, the more energy that is required to melt a substance (higher melting point) • the same concepts applied to boiling points • the stronger the interactions between molecules, the higher the melting and boiling points High Boiling Point

Intermediate Boiling Point

Low Boiling Point

Ionic Compounds

Molecular Compounds (Polar)

Non-Polar Compounds

• Cesium Chloride → CsCl

• Ethanol → C2H5OH

• Hydrogen → H2

• Magnesium Oxide → MgO

• Ammonium → NH4

• Nitrogen → N2

• Sodium Chloride → NaCl

• Hydrogen Chloride → HCl

• Methane → CH4

Ionic Compounds

Polar Molecular Compounds

Non-Polar Molecular Compounds

Strong interactions between molecules

Intermediate interactions between molecules

Weak interactions between molecules

Solubility: • substance will dissolve in a solvent that has similar properties (with respect to polarity) • “like dissolves like” • polar substances will dissolve in polar solutes (like water) • non-polar substances will not dissolve in water (like oil)

Electrical Conductivity: • electric current can flow through substances if electrons or ions are able to flow • metals - as solids, electrons can move between atoms (therefore metals conduct electricity) non-metals - don’t conduct electricity because there atoms do not move • • ionic compounds - as solids, they don’t conduct electricity • ionic compounds - if they are dissolved in water they do conduct electricity • molecular compounds - cannot conduct electricity when solid or when dissolved in water **pure water cannot conduct electricity** **water that is not pure can conduct electricity**

Intermolecular forces: • London forces - the weakest intermolecular force; exists between all molecules but is only significant in non-polar molecule • Dipole-dipole forces - intermediate strength intermolecular force; exists between polar molecules • Hydrogen bonds (or H-bonds) - strongest intermolecular; force exists polar molecules that have H (with a partial positive charge) and O,F, or N (with a partial negative charge). (not a bond)...