Chem Atomic Structure & Chemical Bonding 2009
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Atomic Structure @ @ @ @ @ @ @ @ Neutral atom: no. of electrons = no. of protons Negative ion (anion): no. of electrons > no. of protons Positive ion (cation): no. of electrons < no. of protons Atomic no. = No. of protons No. of protons + No. of neutrons = Mass no. No. of protons = No. of electrons (in atoms) Atoms and Ions have different no.s of electrons Nucleons do not change during ion formation (only electrons) Isotopes @ @ @ @ Same number of protons (same element) Different mass numbers—change in no. of neutrons Same chemical properties Different physical properties (more neutrons, greater mass) First ionization energy @ @ @ Amount of energy required to remove one electron from each atom in a mole of gaseous atoms producing one mole of gaseous cations Factors influencing ionization energies @ Size of positive nuclear charge @ As nuclear charge increases, attraction for outermost electrons increases and hence more energy is required to remove an electron @ Size of atom/ion @ As atomic size increases, attraction of positive nucleus from negative electron decreases, hence less energy required to remove an electron @ Screening/shielding effect of inner electrons @ Outermost electrons shielded from attraction of nucleus @ As shielding increases, attraction decreases, hence Ionization Energy decreases Trends @ Down group: I.E. @ Due to increasing atomic size and increasing shielding @ Across period: I.E. @ Due to increasing nuclear charge and decreasing atomic radius Chemical bonding Formed between 2 atoms using their valence electrons, through TRANSFER or SHARING @ 3 different kinds of bonds @ Ionic @ Covalent @ Metallic Octet Rule @ Eight electrons in their valence shells, completely filled @ Configuration of a noble, unreactive gas @ Gives compound stability @ *Exceptions @ First few electrons (eg Lithium) which form configuration of Helium (2 valence) @ Octet Expanders (eg Sulpur in SO3) which can accommodate more than 8 valence electrons Bonds @ Metallic (positive metal ions to sea of electrons) @ Ionic (positive metal ions to negative non-metal ions) @ Covalent (non-metal ion to non-metal ion) @ @ Between atoms/ions are STRONG Between molecules are WEAKER (eg Hydrogen bonding, Van der Waals’ forces) Ionic @ @ Metal and Non-metal Arises due to electrostatic attraction of positive and negative ions @ Atom can lose electrons positively-charged ion (Cation) @ @ Atom can gain electrons negatively-charged ion (Anion) Dot & Cross Diagrams @ Legend @ Cation does not need valence electrons (only symbol, bracket, charge) @ Anion needs electrons (to show the transfer) When an ionic compound is formed, the valence electrons lost by the metallic elements will be transferred to the outermost shell of the atoms of the non-metallic elements @ Covalent @ @ @ @ @ @ @ @ @ 2 non-metallic elements Sharing of valence electrons Type of bond @ Polar (unequally shared, eg HCl) @ Non-polar (equally shared, eg Cl2) Single bond (Cl-Cl) Double bond (O=O) Triple bond (NN) Dative Bonds Covalent bond in which both electrons come from one atom, which are then shared with another atom To form a dative bond, the donor group must have a lone pair of electrons (electrons that are not involved in bonding) in its outermost shell while the accepter group must have vacant orbital(s) in its outer shell, which are able to accommodate the lone pair of electrons from the donor. Metallic @ @ Formed between cations of a metal and its freely moving delocalized electrons Atoms are held together by the electrical interaction between the positively-charged ions and negativelycharged electrons Electro negativity @ @ @ @ @ @ Measure of the tendency of an atom to attract a bonding pair of electrons When atoms of elements with different electronegativities are bonded together, the more electronegative element attracts the electrons to itself When electrons are not shared equally, the resulting covalent bond = Polar Bond, and it produces a Dipole Moment, represented by arrows towards the more electronegative atom Metals have lower electronegativities Non-metals have higher electronegativities Types of bonds formed depend on difference in electronegativity: -< 1.5: covalent bond ->1.5: ionic bond Dipole Moment @ Physical property that determines the asymmetry of a charge distribution Polar and Non-polar Difference in electronegativity for atoms in bond >0.4 = Polar Difference in electronegativity for atoms in bond <0.4 = Non-Polar If there are no lone pairs on the central atom, and if all the bonds in the central atom are the same, the molecule is non-polar If the central atom has at least one polar bond and if the groups bonded to the central atom are not all identical, the molecule is probably polar Intermolecular Forces Van der Waals’ Forces @ @ @ @ All atoms and molecules experience Van der Waals’ forces Attractive forces arising from fluctuations in the electron distribution within atoms or molecules At any one instant, the random motion of electrons within an atom/molecule may cause the electrons to be clustered more at one end of the particle, giving that end a very small partial negative charge, creating an instantaneous dipole Van der Waals’ forces exist between all molecules, but are the only forces that exist between non-polar molecules Dipole-Dipole Forces @ @ @ Attractions between the opposite partial charges in the permanent dipoles of polar molecules Dipole-dipole forces exist between all polar molecules in addition to Van der Waals’ forces The greater the difference in electronegativity, the stronger the permanent dipole-dipole attraction Hydrogen Bonds @ @ @ @ @ Electronegativity difference between O, N, and F vs H is so large that these bonds are especially POLAR Attractions between the opposite partial charges are especially strong H-bond < Covalent & Ionic bonds Strength of H-bond > Dipole-Dipole force > Van der Waals’ force Importance @ Structure and property of water and ice @ H-bonding results in water molecules to be further apart in ice than in water @ Solubility in water @ Polar substances are soluble in water @ Ability of other molecules capable of H-bonding increase solubility @ Ion-dipole interaction solubility of ionic substances @ Non-polar substances will not dissolve in water @ Dimerization of carboxylic acids @ DNA Answer format for questions 1) What bonding 2) What structure 3) Amount of energy required to break bonds Possible Answer Template: -As _______ is non-polar, only van der Waals forces exist between the molecules and hence little energy is required to overcome the weak intermolecular forces. -________ is polar. Other than van der Waals forces, dipole-dipole forces exist between the molecules. More energy is required to overcome both the van der Waals forces and the dipole-dipole forces. Chemical structures Simple atomic/molecular structures @ @ @ Individual separate units (atoms/molecules) Simple atomic: Eg Helium Simple molecular: Eg Water (H2O), Methane (CH4), Chlorine (Cl2), Glucose (C6H12O6) Giant structures @ @ @ @ @ @ @ @ @ Many particles joined together by strong bonds into a large network No separate units Particles are atoms/ions Substances with giant structures High MP and Bp Solids at RTP Giant metallic: Eg Iron, Copper Macromolecular (also giant atomic/giant molecular): Eg Diamond, Graphite (carbon), Sand (SiO 2), Plastics (CH2)n, etc Ionic: Eg Sodium chloride (NaCl) Properties @ @ @ Volatility (high or low BP/MP) Electrical conductivity (whether it conducts readily or not) Solubility (whether it dissolves readily in water, a polar solvent, or a non-polar organic solvent) Volatility Ionic @ High MP/BP @ @ Eg NaCl MP = 801C @ Eg Al2O3 MP = 2050C Why? @ Large amount of heat required to break strong electrostatic forces of attraction between positive and negative ions (holding ions together) @ Must overcome bonds holding ions in liquid state @ Lots of energy needed to overcome these bonds @ Applications @ Ionic compounds are used as refractory materials (heat-resistant, with high MPs) @ MgO: Lines inside of furnace @ Al2O3: Used inside spark plugs Simple Molecular @ Low MP/BP @ @ @ Eg Methane MP = -182C @ Eg Sulphur MP = 144C Why? @ Intermolecular forces are weak, so they can be overcome easily (eg Van der Waals’ forces) @ Less heat energy required to overcome forces *Strong covalent bonds between atoms are not broken by heat; molecules remain intact when state changes Macromolecular @ High MP/BP @ @ Eg Silicon MP = 1650C @ Eg Diamond MP = 3700C @ Eg Graphite MP = 3300C Why? @ Atoms in lattice held together by strong covalent bonds @ A lot of heat energy required to overcome bonds Metallic @ Relatively high MP/BP @ @ @ Eg Aluminium MP = 660C @ Eg Iron MP = 1535C @ Eg Copper MP = 1083C Why? @ Electrostatic forces of attraction between positive ions and “sea” of delocalized electrons are strong @ A lot of heat energy required to overcome forces *Except Group I metals and mercury (liquid at room temp) have lower MP/BP HOW TO ANSWER? 1. Type of electrostatic force of attraction between positive and negative ions (strong or weak?) 2. Amount of energy required to overcome 3. High or low MP/BP Electrical conductivity For a substance to conduct electricity, it must have mobile charge-carriers (ions/electrons) Ionic @ Conducts only in molten/aqueous state @ Why? @ @ @ Ions can move freely in molten/aqueous state The moving ions (charge-carriers) carry the electric current In solid state, the ions (fixed in solid position in lattice) cannot move, so current cannot flow, cannot conduct electricity Simple Molecular @ Do not conduct in any state @ Why? @ Made of neutral molecules, so there are no mobile charge particles to carry charges and conduct electricity @ *Exceptions: Covalent molecules like HCl, HNO3 and H2SO4 can dissolve in water to form acidic solutions containing free ions (H+) thus allowing an electric current to pass through Macromolecule @ Do not conduct in any state @ Why? @ Made of neutral molecules, so there are no mobile charge particles to carry charges and conduct electricity @ *Exception: Graphite is the only non-metal that is a good conductor, because each carbon atom is bonded to 3 others in the graphite lattice. The 4th valence electron is delocalized within each layer, and acts as a mobile charge carrier to conduct electricity Metallic @ All metals conduct electricity in solid and liquid states @ Metallic lattice: positive ions, “sea” of electrons” – charge carriers Solubility in solvents Ionic @ Most can dissolve in water, a polar solvent @ *Exceptions: carbonates, oxides, hydroxides Simple Molecular @ Most covalent substances don’t dissolve in water, but dissolve in organic non-polar solvents @ Eg Iodine is only partially soluble in water, but dissolves well in tetrachloromethane (organic non-polar solvent) Macromolecular @ Not soluble in any type of solvent Metallic @ Not soluble in any type of solvent @ However, they react with water: not physical, but chemical new substances formed Carbon: Allotropes (different forms of same element, different arrangement and bonding) Diamond Graphite MP (very high) 3700C 3300C Density (high – 3.5g/cm3 2.2g/cm3 sink in water) Appearance Colourless, transparent crystal Black, shiny powder Hardness @ Hardest natural substance known @ Soft @ Used as drill tips for drilling @ Used as a solid lubricant to reduce friction in equipment and in glass cutters engines @ Pencil lead @ Consists of millions of carbon atoms @ Consists of parallel layers of carbon atoms strongly and covalently bonded in @ Although the atoms are covalently bonded tetrahedron units within each layer, the bonds between the @ Making the substance rigid and strong layers are weak and allow the layers to easily @ Diamond is a very hard substance slide against one another @ Graphite is soft and slippery Electrical Does not conduct Conducts conductivity @ Each of the carbon atom’s four @ Each carbon atom is covalently bonded to valence electrons is involved in three other carbon atoms covalent bonding with other carbon @ Leaves each carbon with one valence electron atoms not involved in bonding @ No delocalized electrons to move @ Electron become delocalized, can move freely through the structure to conduct among layers of carbon atoms, conducting electricity electricity
