Periodic Trends Interpret the periodic table to describe
Periodic Trends
• Interpret the periodic table to describe an element’s atomic makeup • Describe trends found in the periodic table with respect to atomic size, ionization energy, and electronegativity.
Periodic Law • Developed by Dmitri Mendeleev: – Originally, he arranged the elements in order of increase atomic mass • For the most part, was able to predict the properties of elements that had not yet been discovered – When arranged by increasing atomic number, the chemical elements display a regular and repeating pattern of chemical and physical properties.
Properties of Elements There is another way of classifying element: – Metals – left side of periodic table • lustrous (shiny), malleable, ductile, and are good conductors of heat and electricity • Solids at room temp except. . . – Nonmetals – right side of periodic table • dull, brittle, nonconductors (insulators) • solids & gases at room temp except. . . – Metalloids or Semi-Conductors – along stair step line • have characteristics of both metals and nonmetals • They are shiny but brittle
Periodic Trends 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.
Effective Nuclear Charge What keeps electrons from simply flying off into space? – the pull that an electron “feels” from the nucleus. – The closer an electron is to the nucleus, the more pull it feels. • As effective nuclear charge increases, the electron cloud is pulled in tighter (making the atom smaller).
Effective Nuclear Charge
Effective Nuclear Charge
Shielding As more energy levels are added to atoms – the inner layers of electrons shield the outer electrons from the pull of the nucleus. – The effective nuclear charge (ENC) on those outer electrons is less • so the outer electrons are less tightly held
Shielding
Shielding
ATOMIC RADIUS the size of the atom • Radius is the distance from the center of the nucleus to the “edge” of the electron cloud – usually measured in angstroms (Å) • 1 angstrom = 1 x 10 -10 m • The trend for atomic radius in a vertical column is to go from smaller at the top to larger at the bottom – Why? • With each step down the group, an entirely new energy level is added, making the atoms larger with each step.
ATOMIC RADIUS • The trend across a horizontal period is less obvious • What happens to atomic structure as we step from left to right? 1. Each step adds a proton and an electron (and 1 or 2 neutrons). 2. Electrons are added to existing PELs or sublevels. • 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 electron cloud closer, making atoms smaller as we move from left to right across a period.
ATOMIC RADIUS In BLACK draw the following arrows to indicate INCREASE in Atomic Radius
ATOMIC RADIUS
Examples – Atomic Radius Rank the following in order of increasing atomic radius: 1. Cl, Mg, Si Cl, Si, Mg 2. Br, F, Cl, Br 3. F, Ca, Al F, Al, Ca 4. Sr, Br, P P, Br, Sr
IONIZATION ENERGY energy required to remove an electron – Measured in kilojoules (k. J) – If an electron is easy to remove, then a LOWER amount of ionization energy is needed – energy is added to the atom to remove the electron –So it is an endothermic process
IONIZATION ENERGY The larger the atom is, the easier its electrons are to remove. –Ionization Energy and Atomic Radius are inversely related. • As atomic radius increase, ionization energy decreases.
IONIZATION ENERGY Consider: Be and F – Which electron is easier and more likely to be removed? Be gives up electrons more readily than F – Meaning, it is EASIER to remove the electron from Be – Therefore, Be has a LOWER ionization energy than F
IONIZATION ENERGY • What is the overall trend of Ionization Energy as you move from LEFT to RIGHT?
IONIZATION ENERGY Increases from Left to right WHY? • # of e- and p+ increase • valence e- are closer to nucleus bc the force of attraction is greater • making it harder to remove an electron
IONIZATION ENERGY In RED draw the following arrows to indicate INCREASE in Ionization Energy
Examples – Ionization Energy Rank the following in order of increasing ionization energy: 5. Cl, Mg, Si, Cl 6. Br, F, Cl Br, Cl, F 7. F, Ca, Al, F 8. Sr, Ar, P Sr, P, Ar
ELECTRONEGATIVITY a measure of an atom’s attraction bonding electrons – arbitrary scale that ranges from 0 to 4 • unit: Paulings • Metals are e- givers • low electronegativities • Nonmetals are e- receivers • high electronegativities • What about the noble gases? • Low to NO electronegativity because they do not form
The Octet Rule • The “goal” of most atoms (except H, Li and Be) is to have an octet or group of 8 electrons in their valence energy level. – accomplished by either giving electrons away or taking them • Metals loses e- (to become positive) • Nonmetals gain e- from other atoms (to become negative) • Atoms that have gained or lost electrons are called ions.
ELECTRONEGATIVITY In GREEN draw the following arrows to indicate INCREASE in Electronegativity
ELECTRONEGATIVITY
Examples – Electronegativity Rank the following in order of increasing electronegativity: 9. Cl, Mg, Si, Cl 10. Br, F, Cl Br, Cl, F 11. F, O, Al Al, O, F 12. Sr, Ar, P Ar, Sr, P
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