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Home , Octet rule

LS50 2015

INTRO TO 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 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 for an • Understand ; 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 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 § tend to lose electrons to become cations § tend to become isoelectric with a (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

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§ tend to become isoelectric with a noble gas § tend to be dull, brittle, poor conductors of heat and electricity § three main types: noble gases, , other non-metals o § 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 § Totally donate them or steal them from another atom à ionic bond § Nothing à no bond o The 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 = 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 (row)

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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 § • 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) § • 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 (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)

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§ Include 2 post-transition metals (Uup, Bi), 2 metalloids (As, Sb) and 2 (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 (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 “” 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. BF3, BeCl2, BCl2 § Valence shell expansion • This tends to happen for third period elements (Na, Mg, Al, Si, P, S. Cl) • Their d orbitals are energetically close enough to be available to participate in bonding • >8 electrons can surround one of these elements in a molecule and it can still form a stable bond • e.g. ClF3, PCl5 § Odd number of valence electrons = unpaired valence electrons • e.g. ClO2 • these are free radicals: very reactive because of their

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Table. Characteristics of Groups (columns) of the Periodic Table

Example Element (main H Be B C N O F He group only) CAS Group # IA IIA IIIB IVB VB VIB VIIB VIII VIII VIII IB IIB IIIA IVA VA VIA VIIA VIIIA (older) IUPAC Group # 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 (modern) # valence electrons 1 2 Transition metals: No good way to predict # valence electrons for these. 3 4 5 6 7 8 (0) Location of Outermost shells can be in their inner shells as well as outer Outermost shells valence electrons Element family Scandium Titanium Vanadium Chromium Manganese Iron Cobalt Nickel Copper Zinc Fluorine name Alkaline Trivial Alkali Coinage Volatile Noble earth Icosagens Crystallogens Halogens name metals metals metals gases metals

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