Write the noble gas notation for the following elements: Germanium (Ge): #32 Niobium (Nb): #41 Osmium (Os): #76

Find the noble gas notation of the following elements: Germanium (Ge): [Ar] 4s2 3d10 4p2 Niobium (Nb): [Kr] 5s2 4d3 Osmium (Os): [Xe] 6s2 4f14 5d6

Orbital Diagrams, Valence Electrons, Lewis Electron Dot Structures and the Periodic Table

Orbital Notation (Diagrams) Since we already know how to do electron configurations of different elements, we can draw what they look like in an orbital diagram All we need to know is the element, how many orbitals it contains and how many electrons can fit within each orbital level! Before we can start, we must follow some rules…

Pauli Exclusion Principle No two electrons in the same atom can have the same set of four quantum numbers This is due to the opposite spins of the electrons within the orbitals Aufbau Principle Orbitals of lowest energy are filled first Hund’s Rule Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron

1s1 Let’s look at Hydrogen’s electron configuration for example: The “s” sublevel of our notation means that we will have one orbital (represented by a circle) that can hold a maximum of 2 electrons. Since we know that the “s” orbital can hold two electrons and we only have one electron to put inside of the circle (Hydrogen is atomic number 1), we just put one arrow (representing an electron) inside of the circle. The arrows that are placed within the circles represent electrons, which MUST have different spins (directions)! Let’s look at another example… 1s1

What about the orbital diagram for Nitrogen? Electron Configuration: 1s2 2s2 2p3 1s 2s 2p Make sure that you always fill the lower energy level orbitals first. They must be full before you can move to the next energy level.

Orbital Diagrams Using our Nitrogen example: 1s2 2s2 2p3 We can write orbital diagrams as a “step” progression as well. Instead of circles, draw lines representing the orbitals in the energy levels and fill in the electrons just as you would using circles. Using our Nitrogen example: 1s2 2s2 2p3 2p ______ ______ ______ 2s ______ 1s ______

Practice: Orbital Diagrams Draw the orbital diagrams for: Carbon Sodium Phosphorus Argon

Valence Electrons

Periods: Energy Levels (n) Each row in the periodic table is called a “period” The period corresponds to a specific energy level of the atom The top row, Period 1, is closest to the nucleus, the next one down is Period 2, etc…until you end with Period 7. Level 1: s Levels 2 and 3: s,p Levels 4 and 5: s,p,d Levels 6 and 7: s,p,d,f

Groups: Valence Electrons Each column in the Periodic Table is called a “group” Each element in a group has the same number of electrons in their outer energy level (the valence level). The electrons in the outer shell are called “Valence Electrons” Red: Group 1 Orange: Group 2 Yellow: Group 13 Green: Group 1 Sky Blue: Group 15 Baby Blue: Group 16 Dark Blue: Group 17 Purple: Group 18

Valence Electrons Valence electrons are the electrons in the highest occupied energy level of the atom. Valence electrons are the only electrons generally involved in bond formation (which we will talk about in the next unit!)

Bohr Atomic Structures Electron Configuration of Na: 1s22s22p63s1 The first energy level contains 2 electrons. (s orbital…1s2) The second level contains 8 electrons. (2s and 2p orbitals…2s22p6) How many electrons do you see in the outermost level? 3s1… 1electron! This is the Valence number. Sodium has 1 Valence electron.

Electron Dot Structure: Lewis Dot Diagrams A notation showing the valence electrons surrounding the atomic symbol. How many valence electrons in Cl? C?

Lewis Dot Structures Find out which group (column) your element is in. This will tell you the number of valence electrons your element has. You will only draw the valence electrons.

N Lewis Structures 1)Write the element symbol. 2) Nitrogen is in the 5th group, so it has 5 valence electrons. 3) Starting at either the right or left of the element symbol, draw 5 electrons, (dots), around the element symbol. N

The History of the Periodic Table and Trends

In the old days… Many elements were known in the ancient world- copper, gold, silver, lead, etc. For several hundred years, elements were discovered by alchemists Alchemy was the ultimate search for wisdom and immortality.

By 1860, more than 60 elements had been discovered….HOWEVER, There was no consistent organization of the elements. No one was using the same method to determine mass of atoms, or the ratios of atoms in compounds In 1860, Stanislao Cannizzaro of Italy presented a convincing method to measure the mass of atoms, thus creating standard values for atomic mass. Now that there are some common standards…

The Matter of Mendeleev In 1869, Dmitri Mendeleev began to try to arrange the elements. Inspired by solitaire, he started to find patterns in the properties of elements, and arranged the known elements by atomic mass in a Periodic (repeating) Table.

Mendeleev’s Genius Mendeleev recognized there were undiscovered elements. By using his periodic table, he could predict the chemical properties of the undiscovered elements. Years later, Scandium, Gallium, and Germanium were discovered and were characteristic of the properties Mendeleev predicted! *swoon*

Henry Moseley 1911: Henry Moseley (working under the direction of Rutherford) rearranged the Periodic Table to go horizontally, and put the elements in order by atomic number.

Variations on the Periodic Table The Mayan Periodic Table of Elements by Mitch Fincher A Spiral Periodic Table by Prof. Thoedor Benfey

The future of the periodic table?

Periodic Trends Now that the periodic table is organized, what patterns can we find? What does it even mean to be “periodic”?

Thanks, to Moseley, we learn that patterns arise because of PROTONS!!! This led to the development of the Periodic Law: the physical & chemical properties of the elements are periodic functions of their atomic numbers. ** In other words, when the elements are arranged in order of increasing atomic number, elements with similar properties appear at regular intervals.

Pattern: Families Elements in column share similar traits, and are called families: These columns are also called groups.

The Alkali Metals 1 valence electron Highly reactive with water Form ionic compounds Do not occur in nature as pure elements (always in compounds)

Alkaline-Earth Metals Have 2 valence electrons Reactive, but less reactive than alkali metals Are ductile, malleable and have a silvery luster Form alkaline (basic) solutions) when put in water

Transition metals… and inner transition metals Are less reactive than groups 1 and 2. Tend not to react in water. Are malleable and ductile, but still harder than group 1 & 2. Tend to be solids at room temperature. Are good conductors of electricity and heat. **Inner transition metals tend to be radioactive

Nonmetals Poor conductors of heat and electricity Often are found as gases or liquids, sometimes solids.

Halogens Are nonmetals highly reactive with metals- most reactive is fluorine, least reactive is astatine Mostly exist as gases or liquids (except At -solid) Have 7 valence electrons

Noble gases At room temperature, exist as gases. Are completely unreactive Have full s and p orbitals Are odorless, colorless, nonflammable

Metalloids Tend to be solids Have properties similar to both metals and nonmetals Tend to be semiconductors (which means they are useful for technological uses)

Valence Electrons: The outermost s & p electrons

Ions Ionization energy Charged atoms that become charged by losing or gaining electrons Ionization energy Energy necessary to make an ion by removing an electron from a neutral atom

Rule #1 to remember! When an element loses an electron, we can think of it as being given away, which is a good thing or POSITIVE thing to do.

Group 1 Elements H, Li, Na, K, Rb, Cs, Fr Achieve a stable octet (full outer shell) by losing 1 electron, which forms a +1 ion H+ , Li+ , Na+ , etc…

Group 2 Elements Be, Mg, Ca, Sr, Ba, Ra Achieve a stable octet by losing 2 electrons, which forms a +2 ion Be2+, Mg2+, Ca2+, etc…

Rule #2 to remember! When an element gains an electron, we can think of it as it is being stolen from another ion, which is a bad or NEGATIVE thing to do.

Group 7 Elements F, Cl, Br, I, At Achieve a stable octet by stealing (gaining) 1 electron, which forms a -1 ion F-, Cl-, Br- , etc…

Group 6 Elements O, S, Se, Te, Po Achieve a stable octet by stealing (gaining) 2 electrons, which forms a 2- ion O2- , S2- , Se2-, etc…

Periodic Trends

What we will investigate: Atomic size How big the atoms are Ionization energy How much energy to remove an electron Ionic size How big ions are Electronegativity The attraction for the electron in a compound

What we know about the atom: All elements have a unique electron configuration, which determines what their atomic and orbital structures will look like A positive nucleus is present and pulls on electrons The more electrons present in the outermost energy level, the stronger the attraction is towards the nucleus of the atom

Keep in mind… + Group trends- as you go down a group: Increase in energy levels (n) Outermost electrons not as attracted by the nucleus because the outermost energy level is further away +

Atomic Size The electron cloud doesn’t have a definite edge, so where do I start to measure?

Atomic Size } Radius Atomic Radius = half the distance between two nuclei of molecule

Trends in Atomic Size Influenced by two factors: Energy Levels (n) Charge on nucleus- More charge pulls electrons in closer

Group Trends- Atomic Size H Group Trends- Atomic Size Li Na As we go down a group, each atom increases by an energy level of n+1 Increased shielding Atoms become larger as you go down group K Rb

Periodic Trends- Atomic Size As we go across a period, the radius gets smaller Same shielding and energy level within a period Increased nuclear charge pulls outermost electrons closer, creating a smaller, more dense atom Na Mg Al Si P S Cl Ar

For each of the following pairs which atom is larger? Mg, Sr Sr, Sn Ge, Sn Ge, Br, Cr, W Sr Sn Ge W

Driving Force for Ionization Energy Full Energy Levels are very low energy Noble Gases have full orbitals (all outer shells end in p6, with the exception of Helium at s2) Atoms behave in ways to achieve noble gas configuration (stable octet)

Ionization Energy Energy necessary to make an ion by removing an electron from a neutral atom Removing one electron makes a +1 ion The energy required is called the first ionization energy

Ionization Energy The second ionization energy is the energy required to remove a second electron, which is always greater than first IE The 3rd IE is the energy required to remove a third electron, which is greater than 1st or 2nd IE

What determines ionization energy? The greater the nuclear charge, the greater IE Increased shielding (energy levels) decreases IE Filled and half filled orbitals have lower energy, so achieving them is easier. This results in lower IE.

Group Trends- Ionization Energy As you go down a group, IE decreases because of more shielding Outer electron is less attracted to the nucleus, so less energy is required to remove that electron

Periodic trends- Ionization Energy All the atoms in the same period have: Same shielding (energy level) Increasing nuclear charge Ionization energy generally increases from left to right. Exceptions: full and half full orbitals hardly require any energy to obtain electrons

Which of the following pairs has a higher ionization energy? Mg, Na S, O Ca, Ba Cl, I Na, Al Se, Br Mg O Ca Cl Al Br

Ionic Size Cations Metals form cations are positive ions form by losing electrons (giving electrons away) Metals form cations Cations of representative elements have noble gas configuration. (8 electrons in outermost shell)

Ionic size Anions are negative ions form by gaining electrons (stealing electrons from other ions) Nonmetals form anions Anions of representative elements have noble gas configuration. (8 electrons in outermost shell)

Configuration of Ions Ions of representative elements have noble gas configuration Example: Na is 1s22s22p63s1 Forms a 1+ ion: 1s22s22p6 (Same configuration as Neon) Metals form ions with the configuration of the noble gas before them on the periodic table- they lose (give away) electrons

Configuration of Ions Non-metals form ions by gaining (stealing) electrons to achieve noble gas configuration. Example: Fluorine is 1s22s22p5 Forms a 1- ion: 1s22s22p6(Same configuration as neon) Non-metals end up with the configuration of the noble gas after them in the periodic table.

Electronegativity The tendency for an atom to attract electrons to itself when it is chemically bonded to another element.

Group Trend- Electronegativity The further down a group: More shielding (energy levels) more electrons an atom has Less attraction for electrons (Low electronegativity)

Periodic Trend- Electronegativity Metals (Left side of Periodic Table) Low nuclear charge Low attraction for electrons Low electronegativity Non-metals (Right side of Periodic Table) High nuclear charge Large attraction for electrons High electronegativity Only exclusion is Group 8: Noble gases… they do not bond with other elements at all

Electronegativity Trend In other words, electronegativity increases across a period and decreases down a group

Which of the following pair has a higher electronegativity? Mg, Na Na, Al Cl, I Ca, Ba S, O Se, Br Mg Al Cl Ca O Br