Unit 3- Atomic Structure Isotope Law of definite proportions Mass number Neutron Orbital Proton Valence Wave-mechanical model Atom Atomic mass Atomic mass unit (amu) Atomic number Electron Excited state Ground state

Atomic theory- atoms are the building blocks of matter Law of definite proportions- a chemical compound will always have the same elements in the same proportions Ex: glucose always C6H12O6 Law of conservation of mass- mass cannot be created or destroyed in ordinary physical or chemical changes Mass of products=mass of reactants 89g oxygen + 11g hydrogen = 100g water

Changing views of atomic structure

1st- John Dalton (1803) 5 principles: now, atoms can be divided 1- all matter is composed of atoms that can’t be created, destroyed or divided now, atoms can be divided technology allows us to create and destroy 2- atoms of an element are identical physically and chemically 3- atoms of different elements differ in physical and chemical properties 4- atoms of different elements combine into compounds (in whole number ratios) 5- in chemical reactions atoms are combined, separated and rearranged, NEVER destroyed, created or changed

2nd- JJ Thompson 1897 Plum pudding model- Experiment with cathode ray tube Cathode ray (-) was deflected by magnetic field Particles must have negative charge- discovered electrons Electrons- particles with negative charge, outside the nucleus

3rd- Ernest Rutherford 1909 Gold foil experiment- discovered positive nucleus and that an atom is mostly empty space Nucleus- dense, central part of atom, positively charged Took a very thin piece of gold foil and shot a beam of alpha particles at it (positively charged) Most went right through thin foil however some were deflected to different degrees He hypothesized a dense positively charged center of the atom -nucleus as almost all the mass but barely any volume -if a marble was a nucleus, the atom would be a football stadium

http://myweb.usf.edu/~mhight/goldfoil.html

4th- Bohr (1913) Electrons travel around the nucleus in specific energy levels

Now- Wave-mechanical model Determined that electrons have properties of mass and also waves Difference from Bohr; instead of electrons being in fixed orbits they are in orbitals Region in which an electron of a particular amount of energy is most likely located (high probability)

The Nucleus Consists of: Protons- positively charged Makes up the atomic number of an atom A neutral atom must have an equal number of electrons and protons Mass: 1amu (atomic mass unit) An atomic mass unit is 1/12th the mass of a C-12 atom Neutrons- neutral in charge Helps hold protons together Neutrons + protons = atomic mass of an atom

O Atomic number- # of protons as atom has Also shows the number of electrons Ex: oxygen has an atomic number of 8; so 8 protons and 8 electrons Mass number- # of protons and neutrons in nucleus Calculate neutrons by subtracting atomic number from mass number 16 8 O

Isotopes atoms of the same element that have a different number of neutrons Atomic number is the same Atomic mass is different Springfield isotopes is the minor league baseball team in the Simpsons ***different elements can have the same mass number but can’t have the same atomic number***

Calculating average atomic mass Most elements are mixtures of isotopes On periodic table- mass is the average If you know the percentage of each isotope you can calculate the average Ex: calculate average mass of nitrogen if a typical sample contains 99.63 % nitrogen-14 and 0.37 % nitrogen-15 Step 1- convert percent to decimal 99.63% = .9963 0.37% = .0037 Step 2- multiply decimals by masses .9963 x 14 = 13.95 .0037 x 15 = .056 Step 3- add products together 13.95 + .056 = 14.006

How many atoms make up that mass? Moles are used by scientists to make sense of the vast number of atoms in a sample The amount of a substance 6.022 x 1023 atoms/particles - Avogadro’s number The number of atoms in exactly 12 grams of C-12 Ex: a dozen eggs, not 12 eggs

Converting with moles Molar mass (gram formula mass) -the mass (g) in 1 mole of a substance Ex: Carbon = 12g/mol Sulfur = 32 g/mol What is the mass in grams of 2.34 mol of lead? 2.34mol x (207.2g/mol) = 484.85g Pb How many moles are in 222g of copper? 222g x (1mol/63.55g) = 3.49 mol Cu http://www2.waterforduhs.k12.wi.us/staffweb/Sciencefolder/Mole%20Day/MOLE_DAY_EXPLOD.gif ---link to mole day animation

Converting with molecules How many atoms are in .632 moles of boron? .632 moles x (6.022x1023) = 3.81x1023 B atoms ** How many atoms are in 6g of neon? 1st- how many moles? 6g x (1mol/20.18g) =.297mol Ne 2nd- moles to atoms: .297mol x (6.022x1023atoms/1mol) = 1.79x1023 Ne atoms

Electron configuration Bohr - Electrons travel around the nucleus in specific energy levels Closer to the nucleus- the lower the E level Like rungs on a ladder, the higher up (farther out) the more potential energy you have, can’t also be in between rungs Now– wave-mechanical model- electrons are in a cloud around the nucleus

Background on light Light is energy Using a prism, light can be broken into different colors. Different colors of light have different amounts of energy. A full spectrum of light has many different amounts of energy.

Flame tests Electrons are normally in a ground state (lowest PE), if it gains energy it moves to an excited state (high PE) In it’s excited state it quickly falls back to ground state and releases ENERGY in the process This release of energy emits a certain wavelength of light Measuring an electrons potential energy tells how far from the nucleus it is

When solutions containing metal ions are heated, they impart characteristic colors to a flame when the electrons fall back down to ground state Sodium Potassium Calcium Barium

Placing the Electrons The location of electrons follows certain rules or principles: The Aufbau Principle − electrons fill orbitals starting at the lowest available energy state before filling higher states. Hund’s Rule − unoccupied orbitals of a given energy will be filled before occupied orbitals of the same energy are reused. (like seats on a bus) The Pauli Exclusion Principle − no two electrons in an atom can occupy the same quantum state. (be in the same orbital and have the same spin) An orbital can hold a maximum of two electrons

Patterns The first energy level has only one sublevel, s; the second energy level has two sublevels, s and p; the third energy level has three sublevels, s, p, and, d; and so on. There is 1 s orbital, 3 p orbitals, 5 d orbitals, 7 f orbitals, and so on. There are a maximum of two electrons per orbital. Using this information, can you fill in the table below? Principal Energy Level (n) Number of Orbitals ( ) Electrons per Sublevel Maximum Number of Electrons ( ) s p d f 1 3 Electrons in Each Location 2 − 4 6 8 5 7 g h i n2 5 7 9 11 13 2n2 9 2 6 10 − − − − 18 16 2 6 10 14 − − − 32 2 6 10 14 18 − − (50) 25 36 2 6 10 14 18 22 − (72) 49 2 6 10 14 18 22 26 (98)

An Analogy Each floor is like an energy level. (1-7) Sometimes it is helpful to think of the atom as an apartment house. Each floor is like an energy level. (1-7) Each apartment is like a sublevel. (s, p, d, or f) Higher energy levels are larger, so they have more sublevels The first level (floor) has one sublevel (apartment) − s The second has two − s and p And so on Each bedroom is like an orbital. “s” apartments have one bedroom “p” apartments have three bedrooms “d” apartments have five bedrooms, and “f” apartments have seven bedrooms

Expanding the Analogy Imagine electrons are moving into the apartment complex pictured below: Electrons don’t like to waste energy climbing to apartments on higher floors. Electrons don’t like to waste energy caring for larger apartments. Electrons move into the most energy efficient apartments first. In what order do the apartments fill up? 12 9 11 6 8 10 4 5 7 2 3 1

Order of Filling The electrons are arranged according to the following rules: the number of electrons equals the number of protons (atomic number) electrons occupy orbitals in sequence beginning with those of the lowest energy. in a given sublevel, a second electron is not added to an orbital until each orbital in the sublevel contains one electron. One consequence of the order of filling is that an outer shell never has more than eight electrons. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6

A Sample Electron Configuration Consider the element scandium shown below: What do the rules for the order of filling show about the electron configuration? First 1s fills with 2 electrons leaving 19 Second 2s fills with 2 electrons leaving 17 Next 2p fills with 6 electrons followed by 3s with 2 electrons leaving 9 Then 3p fills with 6 electrons followed by 4s with 2 electrons leaving 1 Finally the remaining electron goes into 3d This gives an electron configuration of 1s22s22p63s23p64s23d1. 44.9559 +3 Sc 21 2-8-9-2 1s 2s 2p 3s 3p 3d 4s

Types of Electron Configurations Sublevel Notation: sublevel notation shows how many electrons are in each sublevel Bohr Notation: Bohr notation shows the number of electrons in each shell or energy level Orbital Notation: Orbital notation shows the electrons and their spin in each orbital 1s22s22p63s23p64s23d1 2 − 8 − 9 − 2 1s ___ ↑↓ 2s 2p 3s 3p 4s ↑ 3d

More on Orbital Notation Drawing the Orbital Notation for Iron (Atomic Number = 26) Step 1: Determine the sublevel notation of iron 1s has room for 2 electrons, leaving 24. 2s has room for 2 electrons, leaving 22. 2p has room for 6 electrons and 3s has room for 2 more, leaving 14. 3p has room for 6 electrons and 4s has room for 2 more, leaving 6. 3d has room for the remaining 6 The electron configuration is 1s22s22p63s23p64s23d6. Step 2: Use the information from the sublevel notation to draw the orbital notation Draw a horizontal line to represent each orbital in a sublevel, and label it. Add one electron at a time to each orbital, represented by an up or down arrow, in the same order as in the sublevel notation. Electrons in an orbital must have opposite spins (arrows in opposite directions). Follow Hund’s rule. Do not begin pairing electrons in an orbital until all the orbitals in a sublevel have an electron 1s 2s 2p 3s 3p 3d 4s ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↓ ↑ ↑ ↑ ↑ 1s ___ 2s 2p 3s 3p 4s 3d

Summary Draw the sublevel notation by following the order of filling. Draw the Bohr notation by adding together all the electrons in the same energy level. Draw the orbital notation by taking the information from the sublevel notation. 1s2 2s2 2p6 3s2 3p6 3d10 4s2 4p6 4d10 4f14 5s2 5p6 5d10 5f14 6s2 6p6 6d10 7s2 7p6 Fe: 1s22s22p63s23p64s23d6 Fe: 2-8-14-2 1s ___ ↑↓ 2s 2p 3s 3p 4s ↑ 3d ↓ Fe: