Periodic Trends Elemental Properties and Patterns The Periodic

Periodic Trends Elemental Properties and Patterns

The Periodic Law • Dimitri Mendeleev was the first scientist to publish an organized periodic table of the known elements.

The Periodic Law • Mendeleev even went out on a limb and predicted the properties of 2 at the time undiscovered elements. • He was very accurate in his predictions, which led the world to accept his ideas about periodicity and a logical periodic table.

Mendeleev’s Prediction

• Mendeleev arranged the elements by increasing atomic mass. • The modern periodic table is in order of increasing atomic number. • Mosley is given credit for arranging by atomic number.

The Periodic Law • When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.

The Periodic Law • Atoms with similar properties appear in groups or families (vertical columns) on the periodic table. • They are similar because they all have the same number of valence (outer shell) electrons, which governs their chemical behavior.

Valence Electrons • Do you remember how to tell the number of valence electrons for elements in the s- and p-blocks? • How many valence electrons will the atoms in the d-block (transition metals) and the fblock (inner transition metals) have? • Most have 2 valence e-, some only have 1.

A Different Type of Grouping • Besides the 4 blocks of the table, there is another way of classifying elements: • Metals • Nonmetals • Metalloids or Semi-metals

Metals, Nonmetals, Metalloids • There is a zig-zag or staircase line that divides the table. • Metals are on the left of the line, in blue. • Nonmetals are on the right of the line, in orange.

Metals, Nonmetals, Metalloids • Elements that border the stair case, shown in purple are the metalloids or semimetals. • There is one important exception. • Aluminum is more metallic than not.

Metals • Metals are lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity. • They are mostly solids at room temp. • What is one exception? • Hg

Nonmetals • Nonmetals are the opposite. • They are dull, brittle, nonconductors (insulators). • Some are solid, but many are gases, and bromine is a liquid.

Metalloids • Metalloids, aka semi-metals are • • • just that. They have characteristics of both metals and nonmetals. They are shiny but brittle. They are semiconductors. What is our most important semiconductor? Si

Reactivity of Alkali Metals • https: //www. youtube. com/watch? v=UO 0 C KJ 0 ubw. M

Good Afternoon! • You will need your notes, reference packet, and last night’s homework

Periodic Trends

Important Concepts outermost • Valence electrons- electrons in the ______ energy level inner • Core electrons- electrons in the ____ energy levels • When an atom loses an electron, which electron is removed? valence

Effective Nuclear Charge (Zeff) • What keeps electrons from simply flying off into space? • Effective nuclear charge is the pull that an electron “feels” from the nucleus. • The closer an electron is to the nucleus, the more pull it feels. (Higher Zeff) • As effective nuclear charge increases, the electron cloud is pulled in tighter.

Shielding Effect • The core electrons in an atom “shield” the valence electrons from the pull of the nucleus • As the shielding increases (more e- between nucleus and valence e-), the effective decreases nuclear charge _____.

Sodium and Potassium

Applications • • Which atom (Na or K) is a larger atom? Why? More energy levels Which atom has more shielding? K Which atom has a higher effective nuclear charge? Na • Which atom would require less energy to remove the valence electron? (easiest to remove) K

Atomic Radius (Size) • The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom of the family. Increases down the group. • Why? • With each step down the family, we add an entirely new energy level to the electron cloud, making the atoms larger.

Atomic Radius • The trend across a horizontal period is less obvious. • What happens to atomic structure as we move from left to right across a period? • Each step adds a proton and an electron (and 1 or 2 neutrons). • Electrons are added to existing energy levels or sublevels.

Atomic Radius • The effect is that the more positive nucleus has a greater pull on the electron cloud. • The nucleus is more positive and the electron cloud is more negative. • The increased attraction pulls the cloud in, making atoms smaller as we move from left to right across a period. • The Zeff increases but the shielding does not change.

Atomic Radius decreases increases

Practice 1. Arrange the following in order of increasing size: A) O, Se, S O, S, Se B) P, Mg, Cl Cl, P, Mg C) S, F, Rb F, S, Rb D) P, Ga, O O, P, Ga

Ionization Energy • This is the second important periodic trend. • If an electron is given enough energy (in the form of a photon) to overcome the effective nuclear charge holding the electron in the cloud, it can leave the atom completely. • The atom has been “ionized” or charged. • The number of protons and electrons is no longer equal.

Ionization Energy • The energy required to remove an electron from an atom is ionization energy. (measured in kilojoules, k. J) • The larger the atom is, the easier its electrons are to remove. • Ionization energy and atomic radius are inversely proportional. • Ionization energy is always endothermic, which means energy is added to the atom to remove the electron.

Ionization Energy • As you move down a group, the ionization energy decreases. It becomes easier to remove the valence electron. • Why? Bigger atom -> less Zeff -> more shielding • As you move across a period, the ionization energy increases. • Why? Smaller atom -> more Zeff -> less shielding

Ionization Energy increases decreases

Exceptions (Honors) *does not vary uniformly across period* • 2 nd period in order of increasing I. E. • Li, Be, C, Why? O, N, F, Ne

Practice 1. Arrange in order of increasing ionization energy: A) Br, Cl, I I, Br, Cl B) Sb, Sr, I Sr, Sb, I C) (Honors) Ge, As, Se Ge, Se, As

Electron Affinity (Honors) - Electronegativity • What does the word ‘affinity’ mean? • Electron affinity is the energy change that occurs when an atom gains an electron (also measured in k. J). • Where ionization energy is always endothermic, electron affinity is usually exothermic, but not always.

Electron Affinity (Honors) • Electron affinity is exothermic if there is an empty or partially empty orbital for an electron to occupy. • If there are no empty spaces, a new orbital or energy level must be created, making the process endothermic. • This is true for the alkaline earth metals and the noble gases.

Electron Affinity increases decreases Do not include the noble gases!

Metallic Character (Activity of Metals) • Which family of metals would be the most reactive? Alkali Metals • Why? • As you move down a group, the activity of metals increases. • As you move across a period, the activity of metals decreases.

Activity of Metals • The activity of metals is based on how easily the metal loses an electron. • Very active metals lose their valence electrons easily.

Activity of Metals decreases increases

Examples 1. What is the most active metal: A) Ca, Be, Sr B) Sc, K, Ca K C) Pd, Cd, Sr

Nonmetallic Character (Activity of Nonmetals) • What family is the most active of the nonmetals? • Why? • As you move down a group, the activity of nonmetals decreases. • As you move from left to right across a period, the activity of nonmetals increases. • What is the most active nonmetal? F

Activity of Nonmetals • The activity of nonmetals is based on how easily an atom will gain an electron. • Very active nonmetals have a strong attraction for electrons.

Activity of Nonmetals increases decreases Do not include the noble gases!

Way To Remember The Trends What trends increase down a group and decrease across a period? decrease increase size, activity of metals

What is the Name of the Snowman? Size Activity of Metals

Good Afternoon! • You will need your homework, note packet, and reference packet

Ionic Radius • Cations are always smaller than the original atom. • The valence electrons are removed during • • ionization. Ca 2+ is ______than Ca. Conversely, anions are always larger than the original atom. Electrons are added to the outer energy level. S 2 - is _______than S.

Cation Formation Effective nuclear charge on remaining electrons increases. Na atom 1 valence electron 11 p+ Valence e- lost in ion formation Result: a smaller sodium cation, Na+ Remaining e- are pulled in closer to the nucleus. Ionic size decreases.

Anion Formation Chlorine atom with 7 valence e 17 p+ One e- is added to the outer shell. Effective nuclear charge is reduced and the e- cloud expands. A chloride ion is produced. It is larger than the original atom.

Examples 1. Arrange from smallest to largest: A) H, H+, H- H+, H, HB) K+, Cl-, S-2, Ca+2, K+, Cl-, S-2 2. Given the following elements, which would be the largest when forming an ion? Br, Se, Rb, Sr Se (Se-2)

Periodic Table Song https: //www. youtube. com/watch? v=Vg. VQKC cfwn. U&list=PLJr. I 7 At. Iq. GQQYegg. VPRZQx. E ILch. YTSc. ZH&index=42

Ionization Energy (Honors) • First Ionization Energy – the energy required to remove 1 valence electron. • Second Ionization Energy – the energy required to remove 2 valence electrons. • Third Ionization Energy – remove 3 electrons

Ionization Energy • Once the valence electrons are removed, it becomes very difficult to remove the electrons.

Ionization Energy

Ionization Energy 1. Which atom would have a large increase between the 1 st and 2 nd ionization energies? A. Sr B. Ga C. Li Li

Ionization Energy 2. Where would there be the biggest increase in ionization energies? A. Ca 2 nd and 3 rd B. Si 4 th and 5 th C. K 1 st and 2 nd