Ch. 2 Atoms Chemistry II Milbank High School

Ch Dalton’s Atomic Theory The Nuclear Atom Isotopes, Atomic Masses, and Nuclear Symbols

Dalton’s Atomic Theory  All matter is made of tiny indivisible particles called atoms.  Atoms of the same element are identical, those of different atoms are different.  Atoms of different elements combine in whole number ratios to form compounds.  Chemical reactions involve the rearrangement of atoms. No new atoms are created or destroyed.

Just How Small Is an Atom? Think of cutting a piece of lead into smaller and smaller pieces How far can it be cut? An atom is the smallest particle of an element that retains the properties of that element Atoms-very small: Fig. 5.2, p. 108 –still observable with proper instruments: Fig. 5.3, page 108

Counting the Pieces Atomic Number = number of protons in the nucleus # of protons determines kind of atom (since all protons are alike!) the same as the number of electrons in the neutral atom. Mass Number = the number of protons + neutrons. These account for most of mass

Symbols Contain the symbol of the element, the mass number and the atomic number. X Mass number Atomic number

Symbols n if an element has an atomic number of 34 and a mass number of 78 what is the –number of protons –number of neutrons –number of electrons –Complete symbol

Symbols n if an element has 91 protons and 140 neutrons what is the –Atomic number –Mass number –number of electrons –Complete symbol

Isotopes Atoms of the same element can have different numbers of neutrons. different mass numbers. called isotopes.

Naming Isotopes We can also put the mass number after the name of the element. carbon- 12 carbon -14 uranium-235

Atomic Mass How heavy is an atom of oxygen? –There are different kinds of oxygen atoms. More concerned with average atomic mass. Based on abundance of each element in nature. Don’t use grams because the numbers would be too small.

Measuring Atomic Mass Unit is the Atomic Mass Unit (amu) One twelfth the mass of a carbon-12 atom. Each isotope has its own atomic mass, thus we determine the average from percent abundance.

Atomic Mass Magnesium has three isotopes % magnesium 24 with a mass of amu, 10.00% magnesium 25 with a mass of amu, and the rest magnesium 25 with a mass of amu. What is the atomic mass of magnesium? If not told otherwise, the mass of the isotope is the mass number in amu

Ch The Bohr Model of the Atom Electronic Configurations The Aufbau Principal

Ernest Rutherford’s Model Discovered dense positive piece at the center of the atom- nucleus Electrons would surround it Mostly empty space “Nuclear model”

Niels Bohr’s Model He had a question: Why don’t the electrons fall into the nucleus? Move like planets around the sun. In circular orbits at different levels. Amounts of energy separate one level from another. “Planetary model”

Bohr’s planetary model Energy level of an electron analogous to the rungs of a ladder electron cannot exist between energy levels, just like you can’t stand between rungs on ladder Quantum of energy required to move to the next highest level

The Quantum Mechanical Model Energy is quantized. It comes in chunks. A quanta is the amount of energy needed to move from one energy level to another. Since the energy of an atom is never “in between” there must be a quantum leap in energy. Erwin Schrodinger derived an equation that described the energy and position of the electrons in an atom

Things that are very small behave differently from things big enough to see. The quantum mechanical model is a mathematical solution It is not like anything you can see. The Quantum Mechanical Model

Has energy levels for electrons. Orbits are not circular. It can only tell us the probability of finding an electron a certain distance from the nucleus. The Quantum Mechanical Model

The atom is found inside a blurry “electron cloud” A area where there is a chance of finding an electron. Draw a line at 90 % Think of fan blades The Quantum Mechanical Model

Atomic Orbitals Principal Quantum Number (n) = the energy level of the electron. Within each energy level, the complex math of Schrodinger’s equation describes several shapes. These are called atomic orbitals - regions where there is a high probability of finding an electron. Sublevels- like theater seats arranged in sections

Summary s p d f # of shapes Max electrons Starts at energy level

By Energy Level First Energy Level only s orbital only 2 electrons 1s 2 Second Energy Level s and p orbitals are available 2 in s, 6 in p 2s 2 2p 6 8 total electrons

By Energy Level Third energy level s, p, and d orbitals 2 in s, 6 in p, and 10 in d 3s 2 3p 6 3d total electrons Fourth energy level s,p,d, and f orbitals 2 in s, 6 in p, 10 in d, ahd 14 in f 4s 2 4p 6 4d 10 4f total electrons

By Energy Level Any more than the fourth and not all the orbitals will fill up. You simply run out of electrons The orbitals do not fill up in a neat order. The energy levels overlap Lowest energy fill first.

Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 56 electrons

Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 88 electrons

Fill from the bottom up following the arrows 1s 2s 2p 3s 3p 3d 4s 4p 4d 4f 5s 5p 5d 5f 6s 6p 6d 6f 7s 7p 7d 7f 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d 10 5p 6 6s 2 4f 14 5d 10 6p 6 7s 2 5f 14 6d 10 7p electrons

Write these electron configurations Titanium - 22 electrons –1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 2 Vanadium - 23 electrons –1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 3 Chromium - 24 electrons –1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 4 expected –But this is wrong!!

Chromium is actually: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 5 Why? This gives us two half filled orbitals. Slightly lower in energy. The same principal applies to copper.

Copper’s electron configuration Copper has 29 electrons so we expect: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 9 But the actual configuration is: 1s 2 2s 2 2p 6 3s 2 3p 6 4s 1 3d 10 This gives one filled orbital and one half filled orbital. Remember these exceptions: d 4, d 9

Ch Mendeleev’s Periodic Table Electronic Structure and the Modern Periodic Table Valence-Shell Electrons Periodic Table Properties Ionization Energy

Development of the Periodic Table mid-1800s, about 70 elements Dmitri Mendeleev – Russian chemist Arranged elements in order of increasing atomic mass Thus, the first “Periodic Table”

Mendeleev Left blanks for undiscovered elements When discovered, good prediction Problems? –Co and Ni; Ar and K; Te and I

Periodic table Horizontal rows = periods –There are 7 periods Periodic law: Vertical column = group (or family) –Similar physical & chemical prop. –Identified by number & letter

Areas of the periodic table Group A elements = representative elements –Wide range of phys & chem prop. –Metals: electrical conductors, have luster, ductile, malleable

Metals Group IA – alkali metals Group 2A – alkaline earth metals Transition metals and Inner transition metals – Group B All metals are solids at room temperature, except _____.

Nonmetals Nonmetals: generally nonlustrous, poor conductors of electricity –Some gases (O, N, Cl); some are brittle solids (S); one is a fuming dark red liquid (Br) Group 7A – halogens Group 0 – noble gases

Division between metal & nonmetal Heavy, stair-step line Metalloids border the line –Properties intermediate between metals and nonmetals Learn the general behavior and trends of the elements, instead of memorizing each element property

Valence-Shell Electrons The electrons in the outermost shell of the ground-state atoms of the element The elements in a vertical group of the periodic table have the same number of electrons in the valence shell of their atoms Figure 2.14 The period number is the same as the number of the principal shell of the valence-shell electrons

Atomic Radii The distance between the nuclei of two atoms Atomic radii increase from top to bottom within a group of the periodic table Atomic radii of the A-group elements decrease from left to right in a period of the periodic table

Ionization Energy The energy required to remove an electron from a ground-state atom (or ion) in the gaseous state Decrease down a group in the periodic table Increase in going from left to right through a general period

Electron Affinity The energy change that occurs when an electron is added to an atom in the gaseous state Increase as you move right on the periodic table Increase as you move up the periodic table