Aim: How are the elements grouped on the periodic table? List two trends that you know about the periodic table.

Properties of Elements and Trends Aim: How are the elements grouped on the periodic table? The Periodic Table Properties of Elements and Trends

Aim: How are the elements grouped on the periodic table? Mendeleev The original periodic table was created by Dmitri Mendeleev. He organized the table by atomic mass. He predicted existence of 3 missing elements (Ga, Sc, Ge)

Aim: How are the elements grouped on the periodic table? Modern Periodic Table The modern periodic table is organized by atomic number. Henry Mosely

Aim: How are the elements grouped on the periodic table? Vertical columns  called groups or families Horizontal rows  called periods also represent the shells or energy levels The basis of periodic law is repeating patterns.

Major Categories of Elements Aim: How are the elements grouped on the periodic table? Major Categories of Elements 1) metals 2) non-metals 3) metalloids Metals are found to the left of the “staircase” Non-metals are found to the right of “staircase” Metalloids are found “on” the staircase

Aim: How are the elements grouped on the periodic table? Properties of Metals Solid at room temperature (except Hg) Shiny, Have luster Malleable Ductile Generally have high melting points Good conductor of heat and electricity Have metallic bonds

Properties of Non-metals Aim: How are the elements grouped on the periodic table? Properties of Non-metals Can be solid, liquid or gas at room temp S  examples C, S, I2 L  Br2 G  O2, Ne, F2 Allotropes- different molecular forms of the same element (different properties, too) O2 and O3 Dull, lack luster Brittle Poor conductors

Properties of Metalloids Aim: How are the elements grouped on the periodic table? Properties of Metalloids Solid at room temperature Semi-conductors Have varying properties that can be metallic or non-metallic

Aim: How are the elements grouped on the periodic table? Metals tend to lose electrons to form a full valence shell and + ions. Non-metals tend to gain electrons to form a full valence shell and – ions.

Groups on the Periodic Table Aim: How are the elements grouped on the periodic table? Groups on the Periodic Table In general, elements found within the same group have similar chemical and physical properties. This is mainly due to the fact that they have the same number of valence electrons.

Alkali Metals (Group I) Aim: How are the elements grouped on the periodic table? Alkali Metals (Group I) Metals Have one valence electron Tend to lose that one electron to become a more stable ion. Rarely found in natural state.

Lithium(Li), Sodium(Na), and Potassium(K) Aim: How are the elements grouped on the periodic table? Lithium(Li), Sodium(Na), and Potassium(K) Three examples of alkali metals VERY reactive metals Exist as solids at room temperature, but are soft Silvery-white or grayish in color

Alkaline Earth Metals (Group II) Aim: How are the elements grouped on the periodic table? Metals Have two valence electrons Tend to lose the two electrons to become a more stable ion Alkaline Earth Metals (Group II)

Magnesium(Mg), Calcium(Ca) and Strontium(Sr) Aim: How are the elements grouped on the periodic table? Examples of alkaline earth metals Exist as solids at room temperature Silvery white or grayish white in color Fairly reactive metals Magnesium(Mg), Calcium(Ca) and Strontium(Sr)

Aim: What properties of elements represent periodicity 1) Which list includes elements with the most similar chemical properties? a. Br, Ga, Hg b. Cr, Pb, Xe c. O, S, Se d. N, O, F 2) Which element has the highest melting point (Hint: See Reference Table)? a. tantalum b. rhenium c. osmium d. hafnium

Aim: What properties of elements represent periodicity Have 7 valence electrons Non-metals (F, Cl, Br, I) F, Cl (gases), Br (liquid), I (solid) Tend to gain one electron to form a more stable ion Extremely reactive Halogens Group 17

Aim: What properties of elements represent periodicity Very stable because of filled valence shell. Do not react in the presence of other elements (relatively inert) Xe can react with Fluorine (more later) Noble Gases Group 18

Aim: What properties of elements represent periodicity Also known as “heavy metals” Tend to be very dense Have varying reactivity Have multiple oxidation states (can form more than one ion) Have colorful ions (in solution) Transition Metals

Lanthanide and Actinide Series Aim: What properties of elements represent periodicity Rare earth metals Above 92 they are man-made and are unstable Lanthanide and Actinide Series

Periodic Trends Aim: What properties of elements represent periodicity Trends occur in the periodic table in the following areas: -reactivity -metallic/non-metallic character -atomic radius* -ionization energy* -ionic radius -electronegativity* * Found on Table S

Aim: What properties of elements represent periodicity

bigger loser = more reactive Reactivity (metals) Atoms react when they gain, lose or share electrons (bonding). Metals – lose e-, bigger loser = more reactive Across a period, metal reactivity decreases. Down a group, metal reactivity increases.

Metallic Character Remember best metal is Francium (lower left) Metallic character is relative to Fr. Across a period (away from Fr) metallic character decreases. Down a group (towards Fr) metallic character increases. Back

Reactivity (non metals) Atoms react when they gain, lose or share electrons (bonding). non metals – gain e-, Better gainer = more reactive Across a period, non metal reactivity increases. Down a group, non metal reactivity decreases.

Non metallic Character Remember best non metal is Fluorine (upper right) Non metallic character is relative to F. Across a period (towards F) non metallic character increases. Down a group (away from F) non metallic character decreases. Back

Group Trends in Atomic Radius As you move down a group in the periodic table, atomic radius generally increases. This is due to the fact that as you move down a column the number of shells increases making the atom bigger.

Period Trends in Atomic Radius As one moves across a period of the Periodic Table, atomic radius generally decreases. This is due to the fact that each successive element has one more proton that increases the nuclear charge. This positive charge pulls on the outermost shell and makes the atom smaller. Back

Ionization Energy The energy required to remove an electron from an atom is known as the FIRST IONIZATION ENERGY. To remove a second electron requires what is called the SECOND IONIZATION ENERGY. This is due to the fact that as one moves down a group the atomic size increases and outermost e- further away from the nucleus (the nucleus doesn’t have as much pull so it takes less energy to remove).

Period Trends in Ionization Energy As one moves across a period of the Periodic Table, first ionization energy generally increases. This is due the fact that as atomic number (number of protons) increases, so does nuclear charge. The stronger nuclear charge makes it more difficult to remove an electron. Back

An Analogy: Football Remember, as atomic size increases, the outermost electrons become further away from the nucleus. If a football player is holding the ball very close to himself, it is more difficult to remove it from him. If a football player is holding the ball far away from his body is becomes much easier (less energy) to remove it.

Another Analogy: Football Increasing atomic number and therefore nuclear charge would be analogous to the football player building strength. It then becomes more difficult to get the ball from him (ionization energy increases).

Stop and Practice How would you describe the atomic radius of a period-2 alkaline earth metal with that of a period-4 alkaline earth metal? Answer: The atomic radius of a period-2 alkaline earth metal is smaller than that of a period-4 alkaline earth metal.

More practice Which has a larger atomic radius within the same period: a halogen or an alkali metal? Answer: An alkali metal has a larger atomic radius than a halogen within the same period.

More Practice Which element in group 16 has the highest first ionization energy? In period 2? Answer: Oxygen and neon

Aim: How does atomic radius change when you add or subtract electrons? Warm UP: What is ionization energy? What is electronegativity? What is an ion?

Ion Formation Atoms with low ionization energies (they lose electrons easily) form positive ions (CATIONS). Metals do this. Atoms with high ionization energies tend to gain electrons to form negative ions (ANIONS). Non-metals do this.

Ionic Size (Cations) The gaining or loss of an electron affects the size of atom (that turns into an ion). Positive ions are smaller than the atom they are made from.

Cation Size MINUS ONE ELECTRON When an electron is lost from an atom, the nucleus now has more pull on the outer energy level making the ion SMALLER. MINUS ONE ELECTRON

Ionic Size (Anions) Negative ions are larger than the atom they are made from.

Anion Size PLUS ONE ELECTRON When an electron is gained the nucleus now has less pull on the outer energy level making the ion LARGER. PLUS ONE ELECTRON

Trends in Ionic Radius (Both Cations and Anions) Ionic Radius decreases Ionic Radius increases

Electronegativity Can be defined as an atom’s ability to attract electrons. Expressed in arbitrary units on Table S of your Reference Tables. Non-metals tend to have higher electronegativities because they like to gain electrons. atoms electrons

Trends in Electronegativity In one sentence THE CLOSER AN ATOM IS TO FLUORINE ON THE PERIODIC TABLE, THE HIGHER THE ELECTRONEGATIVITY. ** Fluorine has the highest electronegativity on the Periodic Table (4.0).

Stop and Practice How does the ionic radius of sodium compare with that of cesium? Boron with fluorine? Answer: Sodium is smaller than cesium. Boron is larger than fluorine.

Exit Ticket 1) Which statement describes the general trends in metallic properties as the elements in Period 2 are considered in order of increasing atomic number? a. Metallic properties remains same. b. Metallic properties increase. c. Metallic properties increase and then decrease. d. Metallic properties decrease 2) From which of these atoms in the ground state can a valence electron be removed using the least amount of energy? a. nitrogen b. carbon c. oxygen sd. chlorine