• Elements can be metals, nonmetals, or metalloids

Physical Characteristics of Metal

• Malleable • Ductile• Conductive of electricity • Conductive of heat• Have luster and shine• Very High Melting Points

Elemental Classifications

S(s)

O2 (g)

Elemental Classifications

He (g)

Physical Characteristics of Nonmetals• Most nonmetals are gases• Non conductive of heat and electricity• Nonmetal solids are brittle, powdery• Low melting points

• Certain elements are unstable, and hence, do not commonly exist as individual species, but as diatomic molecules

• These include H, O, N, and all of the halogens (group 17)

H H2 Hydrogen (g)O O2 Oxygen (g)N N2 Nitrogen (g)F, Cl, Br, I F2, Cl2, Br2, I2

Fluorine (g), chlorine (g), bromine (L), and iodine (s)

Chemical Groups and Diatomic Molecules

• Molecules are formed by the atomic bonding. There are two types of bonds: IONIC and COVALENT. Special rules exist for naming molecules of each type.

• Ionic bonds exist between metal cation and nonmetal anions

• To name an ionic compound, you do the following1. Write the name of the metal2. Follow it with the ionic name of the nonmetal

Example• KF Potassium Fluoride• MgCl2 Magnesium Chloride

Nomenclature: Ionic Compounds

• Covalent bonds form between nonmetals

• To name an covalent compound:1. Write the name of the first nonmetal.

For non-unity subscripts, use greekprefixes (shown on right)

2. Follow that with the name of the second nonmetal ending with -ide. Again, include greek prefixes. – Only use mono- for oxygen

containing molecules.

Nomenclature: Covalent Compounds

ExamplesCO carbon monoxide N2S dinitrogen sulfideCO2 carbon dioxide P4Se10 tetraphosphorus decaselenide

Nomenclature: Hydrogen

• Hydrogen is strange. It’s a nonmetal, but tends to react in a manner similar to that the metals in the first column.

1. If hydrogen is listed first, use ionic rules to name the molecule.

Ex. H2S = hydrogen sulfide HCl = hydrogen chloride

2. Hydrogen halides that are aqueous (aq, dissolved in water) become acids and are named as such. We drop –gen and add the suffix “–ic acid”

Ex. HCl (aq) = hydrochloric acid ; HF (aq) = hydrofluoric acid

3. If hydrogen is listed last, it is a hydride anion (H-). The naming is either ionic or covalent depending on the molecule.

Ex. MgH2 = magnesium hydride (ionic)TeH2 = tellerium dihydride (covalent)

Group Work

• Name the following:

1. SrO (s)2. IF3 (g)3. HBr (aq)4. CF4 (g)5. NaH (s)6. HCl (g)

• As you can see from the chemicalequation shown to the left, productstypically exhibit vastly different characteristics the reactants

• Also recall our discussion on the law of conservation of mass. Based on this law, can you find a problem with the equation written shown?

𝐍𝐍𝐍𝐍 𝐬𝐬 + 𝐂𝐂𝐥𝐥𝟐𝟐 𝐠𝐠 → 𝐍𝐍𝐍𝐍𝐂𝐂𝐥𝐥 𝐬𝐬 ?

Chemical Reactions

• Mass can not be created or destroyed. This means that every element contained in the reactants must be accounted for in the product(s)

• There are two chlorine atoms on the reactant side, and only one chlorine atom one the product side. To balance the chlorine atoms, we add a coefficient of 2 to the NaCl(s)

• We have balanced the chlorine atoms, but the sodium atoms are now unbalanced. We add a coefficient of 2 to the Na (s). The reaction is now balanced.

𝐍𝐍𝐍𝐍 𝐬𝐬 + 𝐂𝐂𝐥𝐥𝟐𝟐 𝐠𝐠 → 𝐍𝐍𝐍𝐍𝐂𝐂𝐥𝐥 (𝐬𝐬)

Balanced Reactions

𝐍𝐍𝐍𝐍 𝐬𝐬 + 𝐂𝐂𝐥𝐥𝟐𝟐 𝐠𝐠 → 𝟐𝟐 𝐍𝐍𝐍𝐍𝐂𝐂𝐥𝐥 (𝐬𝐬)

𝟐𝟐 𝐍𝐍𝐍𝐍 𝐬𝐬 + 𝐂𝐂𝐥𝐥𝟐𝟐 𝐠𝐠 → 𝟐𝟐 𝐍𝐍𝐍𝐍𝐂𝐂𝐥𝐥 (𝐬𝐬)

• The balanced equation above says that two Na atoms react with one diatomic chlorine molecule to produce two molecules of NaCl

• A coefficient of 2 means that there are two separate species

• A subscript of 2 indicates two atoms bonded together into a single molecule

• Do not confuse coefficients and subscripts. Do not alter subscripts when balancing.

Na Na Cl Cl

NaClNaCl

Coefficients And Subscripts

𝟐𝟐 𝐍𝐍𝐍𝐍 𝐬𝐬 + 𝐂𝐂𝐥𝐥𝟐𝟐 𝐠𝐠 → 𝟐𝟐 𝐍𝐍𝐍𝐍𝐂𝐂𝐥𝐥 (𝐬𝐬)

• Before carrying out any calculations, it is imperative that you first confirm that a given chemical equation is balanced.

• The rules for balancing a chemical equation are provided below.

1. First, balance those elements that appear only once on each side of the equation

2. Balance the other elements as needed. Pay attention to subscripts. In some instances, using fractional coefficients can help.

3. Include phases.

Tips For Balancing Reactions

Balance the following:

1. C3H8(g) + Oxygen (g) Carbon dioxide (g) + Water (L)

2. Sulfur (s) + Oxygen gas (g) Sulfur trioxide (g)

3. NH4ClO4 (s)

Water (L) + Hydrochloric acid (aq) + Nitrogen (g) + Oxygen (g)

4. Hydrochloric acid (aq) + Zinc (s) ZnCl2 (aq) + Hydrogen (g)

Group Work

• As scientists first began to discover and classify the elements, patterns and similarities were observed in chemical behaviors of certain groups of elements.

• Consider the three metals Li, Na, and K– All 3 metals are soft– All 3 metals are less dense than water– All 3 metals have similar appearance and low melting points– The most interesting feature is that all 3 metals react with

the same elements in a nearly identical manner

• As you see in the periodic table, these elements are all listed in the same group, or vertical column.

Chemical Groups And Periodicity

• Dmitri Mendeleev created the periodic table in in 1869 byarranging the elements from left to right in order ofincreasing atomic number, and vertically according totheir behavior (groups)

• In doing so, he observed repetitive patterns in chemicalbehavior across horizontal rows. These patterns areknown as periodicity, and is described in the next slide.

Chemical Groups And Periodicity

Totally unreactive gas

11Na

12Mg

Highly reactive, highly conductive metal

Less reactive, less conductive metal

14Si

17Cl

18Ar

Slightly conductive metalloid

Highly reactive diatomicelement

Totally unreactive gas

19K

20Be

Highly reactive, highly conductive metal

Less reactive, less conductive metal

22Ge

25Br

26Kr

Highly reactive diatomic element

Decreasing metallic character

Decreasing metallic character

Slightly conductive metalloid

Chemical Groups And Periodicity

• The existence of periodicity proves a very important point:– The number of protons in the nucleus has no effect on

chemical behavior. If it were so, chemical behaviors wouldnever repeat since no two elements have the same atomicnumber.

– The chemical behavior of an element is dictated not by thenumber of electrons, but rather, the arrangement of itselectrons around the nucleus.

Explanation Of Elemental Groups

• A direct indication of the arrangement of electrons about a nucleus is given by the ionization energies of the atom

• Ionization energy (IE) is the work function of a gas atom; it is the energy needed to remove an electron (form a cation) Ionizations are successive. – As you remove one electron, it becomes increasingly difficult to

remove the next because of the increasing attraction between the remaining electrons and the protons in the nucleus

𝑀𝑀 → 𝑀𝑀+ + 𝑒𝑒−

𝑀𝑀+ → 𝑀𝑀2+ + 𝑒𝑒−1st Ionization Energy

2nd Ionization Energy

IE1 < IE2 < …….IEn

Ionization Energy

• We can use photons to ionize gas atoms. By measuring the photon energies required to remove electrons from an element, you can gain an idea of how “willing” an atom is to lose an electron, which tells you:– Relative distance of electrons from the nucleus

• Electrons close to the nucleus are very hard to remove. Electrons further away are much easier.

– Relative energy (reactivity) of the electron

• In the next slide, you will see data from an experiment in which the 1st ionization energies of elements are plotted against atomic number.

Ionization Energy

1st Ionization Energy Shows A Periodic Trend

very difficult to ionize

very easy to ionize

• It is relatively easy to remove electrons from metals. It is harder to remove electrons from nonmetals.– It becomes increasingly

difficult as you move right across the periodic table, and up a group.

• It takes a very large amount of energy to ionize a noble gas, so they are usually non-reactive.

The lower the ionization energy of an element, the

more METALLIC and REACTIVE it is.

Ionization Energy

• The closer an electron is to the nucleus, the harder it would be to pull the electron away. – By carrying out multiple ionizations, we can gain

insight into the arrangement of electrons around the nucleus of the element.

Electron Arrangement (Electronic Structure)

Be4 electrons

Li3 electrons

Single electron that is easily removed (far from nucleus)

Pair of electrons that are hard to remove (close to nucleus)

Pair of electrons that are easily removed (far)

• Let’s take a look at the electron configurations of Lithium (atomic # = 3) and Beryllium (atomic # = 4)

Successive Ionizations

Na11 electrons

Ne10 electrons

Same two tightly bound electrons (close)

Eight electrons of similar attraction (distance) to the nucleus

11th electron enters different “shell” (even farther)

Successive Ionizations

Electrons Reside In “Shells” Of Different Distances From The Nucleus

• From these plots, Niels Bohr derived the Bohr model of the atom. In it, electrons reside in shells that orbit at different distances from the nucleus, as discussed in the “Light” lecture.

• Each shell has a finite number of electrons that it can hold• The electrons furthest from the nucleus are the valence

electrons.

• Each shell holds a maxof 2n2 electrons, where the n=1 shell is the closest to the nucleus.

Na

n=1

n=2

n=3

Same Outer Electron Configuration Along A Group Leads to Similarities in Reactivity

Li

Na

KAll group 1 metals have 1 lone electron in the outermost occupied shell (valence shell). This is why they react the same way!!!

Chemical properties of an element are determined by the outer electron configuration.

Periodicity is Due To Repeating Valence Electron Configurations

Li Be B C N O F Ne

Na Al Si P S Cl ArMg

Noble Gas Configurations

• The inner-most electrons of an element comprise what is known as a noble gas core. – At the close of each shell, you have a noble gas

configuration. Noble gases are chemically inactive because they have completely filled shells.

– Elements will react in a way that allows them to either fill or empty their outer shells.

• Lithium, for example, has a two electron core, which we call a Helium core, and valence electron. Sodium has a 10-electron, Neon core, and one valence electron; and so on.

Failure of the Bohr Model

• The Bohr model is a useful representation of the atom. But, Bohr’s mathematical interpretation fails when an atom has more than 1 electron (but, it is still a convenient model).

• In actuality, the electrons are NOT confined to exact orbits as was believed by Bohr. Instead, the electrons are able to move freely within allowed volumes of space called orbitals. Both models agree that an electron’s energy depends on its distance from the nucleus.

Bohr Model

Orbital Model

• An orbital is defined by 4 quantum numbers

n (principle quantum number)L (azimuthal quantum number)mL (magnetic quantum number)ms (magnetic spin quantum number)

Quantum Numbers

• n = 1, 2, 3…..etc. These numbers correlate to the distance ofan electron from the nucleus, very similar to Bohr’s model.

• n determines the energies of the electrons

• n also determines the orbital size. As n increases, the orbitalbecomes larger and the electron is more likely to be foundfarther from the nucleus

1. The Principle Quantum Number, n

• Remember, an orbital is a volume of space where an electron is allowed to reside within an atom. The boundaries of these volumes can take on certain shapes.

• L is restricted to values of 0, 1….(n-1). Each value of L corresponds to a specific orbital “shape”. We will consider the 4 primary types of orbitals

2. The Angular Momentum Quantum Number, L

Orbital Shapes

The different colors of the regions of the orbitals correspond to different magnetic fields (discussed later).

• Can a 2d orbital exist?– 𝑁𝑁𝑁𝑁. 𝐼𝐼𝐼𝐼 𝑛𝑛 = 2, 𝐿𝐿 𝑐𝑐𝑐𝑐𝑛𝑛 𝑁𝑁𝑛𝑛𝑜𝑜𝑜𝑜 𝑏𝑏𝑒𝑒 0 𝑁𝑁𝑜𝑜 1. 𝑂𝑂𝑛𝑛𝑜𝑜𝑜𝑜 2𝑠𝑠 𝑐𝑐𝑛𝑛𝑎𝑎 2𝑝𝑝 𝑁𝑁𝑜𝑜𝑏𝑏𝑜𝑜𝑜𝑜𝑐𝑐𝑜𝑜𝑠𝑠 𝑒𝑒𝑒𝑒𝑜𝑜𝑠𝑠𝑜𝑜. 𝐿𝐿 ≠ 2.

• Can a 1p orbital exist?– 𝑁𝑁𝑁𝑁. 𝐼𝐼𝐼𝐼 𝑛𝑛 = 1, 𝐿𝐿 𝑐𝑐𝑐𝑐𝑛𝑛 𝑁𝑁𝑛𝑛𝑜𝑜𝑜𝑜 𝑏𝑏𝑒𝑒 0. 𝑂𝑂𝑛𝑛𝑜𝑜𝑜𝑜 1𝑠𝑠 𝑁𝑁𝑜𝑜𝑏𝑏𝑜𝑜𝑜𝑜𝑐𝑐𝑜𝑜𝑠𝑠 𝑒𝑒𝑒𝑒𝑜𝑜𝑠𝑠𝑜𝑜. 𝐿𝐿 ≠ 1

• Can a 4s orbital exist?– 𝑌𝑌𝑒𝑒𝑠𝑠 . 𝐼𝐼𝐼𝐼 𝑛𝑛 = 4, 𝐿𝐿 𝑐𝑐𝑐𝑐𝑛𝑛 𝑏𝑏𝑒𝑒 0, 1, 2 𝑁𝑁𝑜𝑜 3. 𝑆𝑆𝑁𝑁 4𝑠𝑠, 4𝑝𝑝, 4𝑎𝑎 𝑐𝑐𝑛𝑛𝑎𝑎 4𝐼𝐼 𝑁𝑁𝑜𝑜𝑏𝑏𝑜𝑜𝑜𝑜𝑐𝑐𝑜𝑜𝑠𝑠 𝑐𝑐𝑜𝑜𝑜𝑜 𝑒𝑒𝑒𝑒𝑜𝑜𝑠𝑠𝑜𝑜.

Examples

• The 3rd quantum number, mL, relates to the spatial orientation of an orbital. mL can assume all integer values between –L and +L

• Number of possible values of mL gives the number of suborbitals of a given type in a specified “shell”

• Based on the specified rules, we see that, for any shell containing these orbitals, there is one s suborbital, three p suborbitals, five d suborbitals, and seven f suborbitals.

3. The Magnetic Quantum Number, mL

3. The Magnetic Quantum Number, mL

• Electrons spin on an axis. Because electrons are charged, their movement creates a magnetic field. Opposing spins create opposing fields. Like spins strongly repel.

• They can spin in two directions with equal probability, as shown below. We call these orientations “spin up” and “spin down”.

• Every suborbital can hold a max of two electrons, one spin up and one spin down.

ms = + 12

(𝑠𝑠𝑝𝑝𝑜𝑜𝑛𝑛 𝑢𝑢𝑝𝑝),−12

(𝑠𝑠𝑝𝑝𝑜𝑜𝑛𝑛 𝑎𝑎𝑁𝑁𝑑𝑑𝑛𝑛)

4. The Magnetic Spin Number, ms

NO TWO ELECTRONS IN THE SAME ATOM CAN HAVE THE SAME 4 QUANTUM NUMBERS!!!

Quantum numbers:(n, L, mL, ms)1, 0 , 0, + ½ 1, 0, 0, – ½ * Allowed

Quantum numbers: 1, 0, 0, + ½ 1, 0, 0 , + ½ * Forbidden !!

1s 1s

Pauli Exclusion Principle

• Now that we have our 4 quantum numbers established, we can use them as a way of “naming” and distinguishing the electrons in an atom. A fundamental rule of electrons is the Pauli exclusion principle:

• n = 2

• L = 0 (s), 1 (p)

• mL = 0 (L=0)= -1, 0, 1 (L=1)

• ms = +/- ½

S

P

Example: List ALL Possible Sets of Quantum Numbers In the n=2 Shell

Three 2p suborbitals𝑚𝑚𝐿𝐿 = −1 𝑚𝑚𝐿𝐿 = 0 𝑚𝑚𝐿𝐿 = 1

(𝑃𝑃𝑒𝑒) (𝑃𝑃𝑜𝑜) (𝑃𝑃𝑃𝑃)

(𝑠𝑠)One 2s orbital

• The diagram to the left shows the order in which electrons populate specific orbitals. They are listed in order of increasing energy.

• Orbitals having the same n, but different L (like 3s, 3p, 3d) have different energies.

REMEMBER: S-orbitals can hold no more than TWO electrons. P-orbitals can hold no more than SIX, and D-orbitals can hold no more than TEN electrons, F holds up to FOURTEEN.

Electron Configurations

• Write the full electron configurations of N, Cl, and Ca

𝑁𝑁 (7 𝑒𝑒𝑜𝑜𝑒𝑒𝑐𝑐𝑜𝑜𝑜𝑜𝑁𝑁𝑛𝑛𝑠𝑠): 1𝑠𝑠2 2𝑠𝑠2 2𝑝𝑝3

𝐶𝐶𝑜𝑜 (17 𝑒𝑒𝑜𝑜𝑒𝑒𝑐𝑐𝑜𝑜𝑜𝑜𝑁𝑁𝑛𝑛𝑠𝑠): 1𝑠𝑠2 2𝑠𝑠2 2𝑝𝑝6 3𝑠𝑠23𝑝𝑝5

𝐶𝐶𝑐𝑐 (20 𝑒𝑒𝑜𝑜𝑒𝑒𝑐𝑐𝑜𝑜𝑜𝑜𝑁𝑁𝑛𝑛𝑠𝑠): 1𝑠𝑠2 2𝑠𝑠2 2𝑝𝑝6 3𝑠𝑠23𝑝𝑝64𝑠𝑠2

Energy

Example

• If we drew the orbital representations of N based on the configuration in the previous slide, we would obtain:

1s

2s

2p

Energy

For any set of orbitals of the same energy, fill the orbitals one electron at a time with parallel spins. (Hund’s Rule)

Drawing Orbital Representations

• We generally recognize only the valence electrons of an atom. We can simply the configuration of the core electrons by using a noble gas notation.

Noble Gas Configurations

ns1

1

2

3

4

5

6

7

ns2ns2

np1ns2

np2ns2

np3ns2

np4ns2

np5

ns2

np6

ns2 (n-1)dx

• Give the noble gas configurations of:• K• K+

• Cl-

• Zn• Te

Group Examples

• When we fill orbitals in order, we obtain the ground state(lowest energy) configuration of an atom.

• What happens to the electron configuration when we excite an electron?

• The electron will move up to the next available orbital. This is the 1st excited state. The next level up is the 2nd

excited state, and so on.

Excited States

• Ground state Li: 1s2 2s1

• 1st excited state Li: 1s2 2p1

• 2nd excited state Li: 1s2 3s1

1s

2s

2p

3s

Ground state

1st excited state

2nd excited state Energy

Example

• As you know, the d-orbitals hold a max of 10 electrons

• These d-orbitals, when possible, will assume a half-filled, or fully-filled configuration by taking an electron from the ns orbital

• This occurs when a transition metal has 4 or 9 valence d electrons with available s-electrons.

4s

3d

Unfavorable4s

3d

Favorable

[Ar] 4s1 3d5Example: Cr [Ar] 4s2 3d4

Transition Metals