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Periodic Table and Periodic Law

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Presentation on theme: "Periodic Table and Periodic Law"— Presentation transcript:

1 Periodic Table and Periodic Law
Chapter 6

2 Chapter Big Idea Periodic trends in the properties of atoms allow us to predict physical and chemical properties!

3 Section 1: Development of the Modern Periodic Table

4 Essential Questions & Vocabulary
How was the periodic table developed? What are the key features of the periodic table? Vocabulary Period law Inner transition metal Group Lanthanide series Period Actinide series Representative element Nonmetal Transition element Halogen Alkali metal Noble gas Alkaline earth metal metalloid Transitional metal

5 Section 1: Main Idea The periodic table evolved over time as scientists discovered more useful ways to compare and organize the elements.

6 Lavoisier In the 1700s, Lavoisier compiled a list of all the known elements of the time.

7 John Newlands (1864) The 1800s brought large amounts of information and scientists needed a way to organize knowledge about elements. John Newlands proposed an arrangement where elements were ordered by increasing atomic mass. Newlands noticed when the elements were arranged by increasing atomic mass, their properties repeated every eighth element (law of octaves).

8 Meyer and Mendeleev (1869) Meyer and Mendeleev both demonstrated a connection between atomic mass and elemental properties. Mendeleev is given more credit because he published his information first. Arranged elements in order of increasing atomic mass into columns with similar properties. Mendeleev was able to predict the existence and properties of undiscovered elements (blank spaces).

9 Mendeleev’s Periodic Table

10 Moseley A few elements were discovered and the atomic masses of the known elements were not accurately determined. Several elements in Mendeleev’s table not in the correct order. Moseley discovered that each element contain a unique number of protons in their nuclei – atomic number Moseley rearranged the table by increasing atomic number, and resulted in a clear periodic pattern.

11 Moseley’s Periodic Table

12 Periodic Law Periodic repetition of chemical and physical properties of the elements when they are arranged by increasing atomic number Periodic comes from the Greek word “periodos”, meaning way around, circuit


14 The Modern Periodic Table
The modern periodic table contains boxes that contain the element's name, symbol, atomic number, and atomic mass.

15 Modern Period Table Groups - Columns of elements
Periods - Rows of elements Representative Elements - Elements in groups 1,2, and 13–18 possess a wide variety of chemical and physical properties Transition Elements -Elements in groups 3–12 Elements are classified as metals, nonmetals, and metalloids.

16 Metals Metals are elements that are generally shiny when smooth and clean, solid at room temperature, and good conductors of heat and electricity. Alkali metals are all the elements in group 1 except hydrogen, and are very reactive. Alkaline earth metals are in group 2, and are also highly reactive. The transition elements are divided into transition metals and inner transition metals. The two sets of inner transition metals are called the lanthanide series and actinide series and are located at the bottom of the periodic table.

17 Nonmetals & Metalloids
Nonmetals are elements that are generally gases or brittle, dull-looking solids, and poor conductors of heat and electricity. Halogens – elements in Group 17 which are highly reactive Noble Gases – elements in Group 18 gases at room temperature and are extremely unreactive Metalloids, such as silicon and germanium, have physical and chemical properties of both metals and nonmetals.


19 Section 2: Classification of Elements

20 Essential Questions & Vocabulary
Why do elements in the same group have similar properties? Based on their electron configurations, what are the four blocks of the periodic table? Vocabulary Valence electron

21 Section 2: Main Idea Elements are organized into different blocks in the periodic table according to their electron configurations.

22 Electron Configuration
Recall electrons in the highest principal energy level are called valence electrons. Only s & p orbital electrons count as valence electrons All group 1 elements have one valence electron.

23 Valence Electrons - Practice
How many valence electrons are in the following elements? Element # Valence Electrons Na Mg H He Cl Al

24 Valence Electrons – Practice Solutions
How many valence electrons are in the following elements? Element # Valence Electrons Na 1 Mg 2 H He Cl 7 Al 3

25 Lewis Dot Structures Draw the Lewis dot structures for each of those elements Na He Mg Cl H Al

26 Lewis Dot Structures of Representative Elements

27 The s-, p-, d-, and f-Block Elements
The shape of the periodic table becomes clear if it is divided into blocks representing the atom’s energy sublevel being filled with valence electrons.

28 S- block Chemically reactive metals Group #1= alkali metals
Slippery appearance and can be cut with a knife! For real? They all have one valence electron Combine readily with the halogens to form salts Group #2- Alkaline Earth Metals Harder, denser and stronger than group #1 metals Have 2 valence electrons

29 d- block - Transition Metals
Lowest quantum # = 3 Maximum # of electrons = 10 There are exception to the electron configuration rules Some metals may form several different ions They are all metals and good conductors of heat and electricity and have high luster. Properties vary greatly. Some metals are highly reactive Other metals not so much- Au, Pt, Pd

30 p-block – Groups 13-18 Group 13- 3 valence electrons
Group valence electrons etc. etc. Contains metals, non-metals and metalloids. Important group- #17- Halogens Most are gases- most reactive with metals

31 Main Group elements aka Representative Elements
Elements found in the s block and p block Only elements that can be used in Lewis Dot Structures

32 Practice Problem Strontium, which is used to produce red fireworks, has an electron configuration of [Kr]5s2. Without using the periodic table, determine the group, period, and block of strontium.

33 Section 3:Periodic Trends

34 Essential Questions & Vocabulary
What are the period and group trends of different properties? How are period and groups trends in atomic radii related to electron configuration? Vocabulary Energy level of an atom Octet rule Ion electronegativity Ionization energy

35 Section 3: Main Idea Trends among elements in the periodic table include their sizes and their abilities to lose or attract electrons.

36 Atomic Radius Atomic size is a periodic trend influenced by electron configuration. For metals, atomic radius is half the distance between adjacent nuclei in a crystal of the element.

37 Atomic Radius For elements that occur as molecules, the atomic radius is half the distance between nuclei of identical atoms that are chemically bonded together.

38 Atomic Radius Trend

39 Shielding/Screening Electrons have an attraction or pull towards the nucleus of the atom (opposite charges attract) Electrons are also repelled away from the inner electrons (like charges repel) Shielding/ Screening: the attraction of valence (outer-shell) electrons is counterbalanced by the repulsion of the inner-shell electrons. The inner-shell electrons “screen” or “shield” the outer-shell electrons from full attraction

40 Effective Nuclear Charge
Effective nuclear charge is the net positive charge experienced by valence electrons.

41 Atomic Radius Decreases From Left to Right
Why? Increasing positive charge in nucleus Valence electrons are not shielded from the increasing nuclear charge Increasing effective nuclear charge

42 Atomic Radius Increases Down a Group
Why? Additional electron shells make the atom larger. Increases

43 Atomic Radii – Practice I
Rank the following atoms in increasing atomic radius. Carbon Beryllium Fluorine Lithium Fluorine < Carbon < Beryllium < Lithium

44 Ions An ion is an atom or bonded group of atoms with a positive or negative charge. Atoms become charged by either gaining or losing electrons. Cations: atoms lose electrons and become positively charged Anions: atoms gain electrons and become negatively charged

45 Cations are smaller than the neutral atom
The loss of a valence electron can leave an empty outer orbital, resulting in a smaller radius. Electrostatic repulsion decreases allowing the electrons to be pulled closer to the nucleus

46 Anions – Bigger than the neutral atom
Why? The addition of an electron increases electrostatic repulsion.

47 Ionic Radius The ionic radii positive ions (cations) generally decrease from left to right. The ionic radii of negative ions (anions) generally decrease from left to right, beginning with group 15 or 16. Both positive and negative ions increase in size moving down a group.

48 Ionic Radius

49 Ionic Radii – Practice Mg2+ Na+ Ti3+
Arrange the following ions in order of increasing ionic radii: Na+, Ti3+, Mg2+ Mg Na+ Ti3+

50 Atomic & Ionic Radii- Mixed Practice
B C If the figure represents the atoms helium, krypton, and radon, match the letter to the correct atom. If the figure represents a cation, an anion, and a neutral atom from the same period, match the letter to correct term. A – Radon B – Krypton C - Helium A – Anion B – Atom C - Cation

51 Ionization Energy Ionization energy is defined as the energy required to remove an electron from a gaseous atom. First Ionization Energy - the energy required to remove the first electron Increases from left to right across a period. Decreases from top to bottom in a group

52 Second Ionization Energy
Second Ionization Energy – the energy required to remove a second electron Each successive ionization requires more energy, but it is not a steady increase.

53 Ionization Energy WHY? Electrons are attracted to the nucleus. The closer they are to the nucleus, the stronger the attraction and the higher the energy needed in order to remove the electron. Opposite trend of Atomic Radius Smaller radii – higher ionization energy

54 Ionization Energy - Practice
Arrange the following elements in order of decreasing Ionization Energy. Al, Mg, Na, Si Si Al Mg Na

55 Interesting IE Pattern
The ionization at which the large increase in energy occurs is related to the number of valence electrons.

56 Octet Rule Octet rule - states that atoms tend to gain, lose or share electrons in order to acquire a full set of eight valence electrons. The octet rule is useful for predicting what types of ions an element is likely to form.

57 Electronegativity Ability for an atom to attract electrons
When it is chemically combined with another atom. Elements with high electronegativities (nonmetals) often gain electrons to form anions. Elements with low electronegativities (metals) often lose electrons to form cations.

58 Fluorine has the highest electronegativity
Increases from left to right Decreases from top to bottom Opposite trend of Atomic Radius. Smaller radii – higher electronegativity (closer electron can get to the nucleus)

59 Electronegativity Visual
Which visual representation best describes electronegativity? The ability of a nucleus of one atom to attract an electron from another atom in a chemical bond.

60 Electronegativity - Practice
Arrange the following in increasing order of electronegativity: Na, Li, K Ca, Br, Se K Na Li Ca Se Br

61 Review of Periodic Trends

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