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Intro to Chemistry Week: September 14 – 18, 2015

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Intro to Chemistry Week: September 14 – 18, 2015 LS50 2015 INTRO TO CHEMISTRY WEEK: SEPTEMBER 14 – 18, 2015 Learning goals By the end of this week, you should understand the following: • The key findings and interpretations of some critical experiments that led to our current understanding of atomic structure and electron configurations (Lecture 08) • How electron configurations are related to chemical bonds and interactions (Lecture 09) • How chemical bonds are related to assembly and structure of biological molecules (Lectures 09- 10) • How molecular assembly and structure are related to molecular function (Lectures 11-12) • How knowing molecular function facilitates prediction of the evolution of biological molecules (Lecture 12) Lecture 09 – Elements and periodic trends Learning goals By the end of this lecture, you should be able to: • Know how to figure out the number of valence electrons for an atom • Understand electronegativity; use this to predict formation of compounds in general terms • Describe and explain some basic properties of an element based on periodic trends • State the octet rule o name and define the types of bonds that atoms can engage in to satisfy it o recognize the three major types of violations of this rule • Draw Lewis dot structures of atoms, ionic compounds, and covalent compounds Lecture Outline Periodic table • note that atomic number is not the same as atomic weight: o atomic number = number of protons (this is the order the elements are arranged in, and is constant regardless of the number of neutrons) o atomic weight = weighted average of the atomic mass numbers of all isotopes of an element o atomic mass number = (atomic number) + (number of neutrons) = (number of protons) + (number of neutrons) • atoms can have more than one o ionic form (with positive (cation) or negative (anion) charge) § charge is determined by the relative number of protons and electrons: § charge = (number of protons) – (number of electrons) o isotope, i.e. have a different number of neutrons § isotopes have the same atomic number but different atomic mass numbers and different numbers of neutrons • the elements break down into some broad categories (we will refine these later): o metals § tend to lose electrons to become cations § tend to become isoelectric with a noble gas (this means that they achieve the same number of electrons as a noble gas) § tend to be shiny (have luster), malleable, ductile, good conductors of heat and electricity § six main types: alkali metals, alkaline earth metals, lanthanoids, actinoids, transition metals and post-transition metals o non-metals § tend to gain electrons to become anions Extavour Lecture Roadmaps: Chemistry Page 1 of 5 LS50 2015 § tend to become isoelectric with a noble gas § tend to be dull, brittle, poor conductors of heat and electricity § three main types: noble gases, halogens, other non-metals o metalloids § can gain or lose electrons: no good way to predict what these will do, although Al tends to be mostly metallic and B tends to be mostly nonmetallic § they border the amphoteric line on the periodic table • Valence electrons o These are the electrons that can participate in forming chemical bonds! o You can use the periodic table to help you figure out how many valence electrons an atom has (see Table) o Noble gases don’t have any valence electrons because their outermost shells are totally full! That’s why they are inert o Atoms can do different things with their valence electrons: § Contribute them to a shared partnership with another atom à covalent bond § Totally donate them or steal them from another atom à ionic bond § Nothing à no bond o The valence electron number and position determine which one of these things will happen = the chemical properties of the element • The octet rule o (it is an empirical observation that) the atoms in the main group elements bond so that each atom has 8 electrons (or 2 electrons if it is H or He) in its outermost shell o can be satisfied by covalent bonds or ionic bonds § ionic compounds: • form when the difference in electronegativity between two atoms is large • an electron is totally transferred from one atom to another (the one with higher electronegativity: further to the right of the periodic table), so that they both satisfy the octet rule § covalent compounds: • form when the difference in electronegativity between two atoms is small(er) • an electron is shared between two atoms, so that they both satisfy the octet rule • Lewis dot structures of atoms o A way to describe valence electrons in atoms, and bonding in polyatomic systems o One electron pair = one single bond (between two atoms) o Follow these steps to draw the Lewis dot structure for atoms: § Draw the atomic symbol § Represent each valence electron as a dot around it o The location of the dots is not important; they will get moved around when assigning electrons to bonds in compounds o When drawing ions with partially filled d or f subshells (e.g. transition metals), typically we don’t include those subshells in the drawing o For cations or anions that might have up to 8 electrons in their outer shell, typically we write the original valence shell configuration (the shell configuration pre-ionization) • Periodic trends o The elements in a given group (column) have similar properties because their outer electron configurations are similar to each other o There are also trends across a period (row) Extavour Lecture Roadmaps: Chemistry Page 2 of 5 LS50 2015 o As long as we arrange elements according to atomic number (assuming neutral atoms, so that # protons = # electrons), the physical and chemical properties of the elements vary periodically o Four major trends that help us understand how atoms behave in chemical bonds: § Atomic radius (and ionic radius) • Decreases across a period from L to R • Increases down a group § Ionization energy • This is the minimum energy needed to remove the highest energy electron from an atom (creating a cation) • Increases across a period from L to R (gets harder to remove electrons L to R across a period) • Decreases down a group (gets easier to remove electrons going down a group) § Electronegativity • This is the net ability of an atom to take an electron from another atom (i.e. be an oxidizing agent) • Increases across a period from L to R (the elements get better at stealing electrons) • Decreases down a group (the elements get worse at stealing electrons) § Electron Affinity • This is the energy change required to add an electron to a neutral atom (creating an anion) • Increases from bottom left to top right (gets harder to add electrons in this direction) o A bit more detail about characteristics of some of the groups: o Alkali metals (Li, Na, K, Rb, Cs, Fr) § Group IA = Group 1 à one valence electron (in s subshell) § Have the largest atomic radius of their period § React violently with water à H2 § Easily ionize to 1+ by losing an electron from s à empty s § Oxidize (lose electrons) in air o Alkali earth metals (Be, Mg, Ca, Sr, Ba, Ra) § Group IIA = Group 2 à two valence electrons (in s subshell) § Easily ionize to 2+ by losing 2 electrons from s à empty s § Close (filled) subshell § React with water à H2 o Transition metals (Sc-Zn, Y-Cd, Hf-Hg, Rf-Cn) § Groups 3-12: multiple valence electrons in d subshells § May have more than one oxidation group § Reactive with acids § Make coloured compounds § Some groups have some some cool trivial names! See Table o Post-Transition metals (Al, Ga, In, Tl, Uut) § Group IIIA = Group 3A à Three valence electrons (in p subshell) § Can have several oxidation states, but commonly 3+ o Group IVA = Group 4A (C, Si, Ge, Sn, Pb, Fl) § à 4 valence electrons (in p subshell) § Includes 1 nonmetal (C), 2 metalloids (Si, Ge) and 3 post-transition metals (Sn, Pb, Fl) § Oxidation numbers +4 to -4 § They form most covalent compounds o Group VA = Group 5A (N, P, As, Sb, Bi Uup) § à five valence electrons (in p subshell) Extavour Lecture Roadmaps: Chemistry Page 3 of 5 LS50 2015 § Include 2 post-transition metals (Uup, Bi), 2 metalloids (As, Sb) and 2 nonmetals (N, P) § Usually form anions -1, -2 or -3 o Group VIA = Group 6A (O, S, Se, Te, Po, Lv) § à six valence electrons (in p subshell) § Include 3 nonmetals (O, S, Se), 2 metalloids (Te, Po) and 1 post-transition metal (Lv) o Halogens (F, Cl, Br, I, At, Uus) § Group VIIA = Group 7A à seven valence electrons (full p subshell, one unpaired in s subshell) § Form monoatomic gases o Noble gases (He, Ne, Ar, Kr, Xe, Rn, Uuo) § Group VIIIA – Group 8A à eight valence electrons (n level orbitals all filled) § Closed shell § Form monoatomic gases o Lanthanoids (La to Lu) and Actinoids (Ac to Lr) § Have a variable number of valence electrons in the f subshell • Lewis dot structures of ionic compounds • Since electrons are not shared between atoms to achieve the octet rule in this type of bond, just draw the structure appropriate for each ion following the electron transfer • One of the atoms involved might end up with zero dots, and the other with eight dots • If necessary indicate the charge in the top right hand corner of the atom • The strength of the bond depends on the charge magnitude (greater charge à stronger bond) and the ion size (smaller ion à stronger bond) o Bond strength is called lattice energy, measured in kJ/mol • Lewis dot structures of covalent compounds o Single bond = one electron pair shared = bond order 1 o Double bond – two electron pairs shared = bond order 2 o Triple bond = three electron pairs shared = bond order 3 o Quadruple bond = four electron pairs shared = bond order 4 o Fractional bond orders are also possible: see “Resonance” below) o Sometimes the octet rule is violated! There are three main examples of this: § Sub-octet systems = octet-deficient molecules • B and Be are prone to this: they can form stable molecules that don’t obey the octet rule • e.g.
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