Atoms and The Periodic Table

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1 Atoms and The Periodic Table A. Early Models of the Atom 1. The earliest models of the atom came in the 5 th century B.C. when In the 4 th century, B.C., rejected this idea and proposed that earthly matter had no properties itself. Instead, It is important to note that Greek theories were based on. They were essentially and in nature and did little or nothing to suggest a direction for experimental work. During the middle ages, This large body of experimental work contained information on VERSION: March 3, 2000

2 2 CHEMISTRY 11 However, as with the Greeks, 2. In 1808, John Dalton reintroduced the idea of atoms and supported his atomic theory on firm experimental foundation. Dalton s atomic theory states that i) Elements are ii) The atoms making up a particular element are iii) Each chemical compound is iv) Chemical reactions involve the The new compounds are These hypotheses explained three fundamental laws: The Law of Definite Proportions, The Law of Multiple Proportions, and The Law of the Conservation of Mass. THE LAW OF DEFINITE PROPORTIONS (Explained by Dalton s and hypotheses.)

3 UNIT VIII ATOMS AND THE PERIODIC TABLE 3 THE LAW OF MULTIPLE PROPORTIONS (Explained by Dalton s hypotheses.) THE LAW OF THE CONSERVATION OF MASS (Explained by Dalton s hypotheses.) In addition to the atomic theory, Dalton made a huge contribution to chemistry by This allowed 3. In 1897, J.J. Thomson discovered that atoms contained Later, he showed that atoms also contained positively-charged particles. Thomson proposed an arrangement for the positively and negatively charge particles inside an atom which was nicknamed.

4 4 CHEMISTRY 11 THOMSON MODEL OF THE ATOM 4. In 1911, Sir Ernest Rutherford This experiment showed that the atom was mostly but contained a Rutherford postulated that

5 UNIT VIII ATOMS AND THE PERIODIC TABLE 5 RUTHERFORD MODEL OF THE ATOM 5. The planeta ry model of electron behaviour proposed by Rutherford suggested that The problem was that This would eventually 6. In 1913, Niels Bohr came up with an equation that accurately predicted the pattern of energies that can be produced by a hydrogen atoms. In order to derive his equation, Bohr suggested that Since electrons could only exist in these orbits,

6 6 CHEMISTRY 11 BOHR MODEL OF THE ATOM Bohr s model was very successful for hydrogen but it ran into problems because

7 UNIT VIII ATOMS AND THE PERIODIC TABLE 7 B. Atomic Numbers and Atomic Mass 1. Chemical elements differ from one another by H has proton in its nucleus He has protons in its nucleus Cl has protons in its nucleus Conversely, any atom having 1 proton must be hydrogen, any atom having 2 protons must be helium, any atom having 17 protons must be chlorine, 2. The ATOMIC NUMBER of an atom = The ATOMIC NUMBER of an atom = In a neutral atom, Neutral Atom number of = number of When the number of protons and electrons are not equal

8 8 CHEMISTRY 11 Adding negative electrons produce a ion while taking electrons away results in a ion. For Ions number of = The ATOMIC MASS of an atom is number of = atomic atomic (protons) If the values for the atomic mass and atomic number must be shown with the atomic symbol, the following super/subscript symbol is used. This symbol is also written as 23 Na (since all sodium atoms have an atomic number of 11) or as Na 23 (avoids having to write atomic mass as superscript)

9 UNIT VIII ATOMS AND THE PERIODIC TABLE 9 EXAMPLE VIII.1 Problem DETERMINING NUMBER OF SUBATOMIC PARTICLES How many protons, electrons, and neutrons do Fe, Al 3+, N 3- and 235 U 2+ contain? Solution protons = since Fe is neutral, electrons = protons = neutrons = = protons = since N 3-, add 3 electrons, electrons = + = neutrons = = protons = 13 since Al 3+, subtract 3 electrons, electrons = = neutrons = = protons = since 235 U 2+, subtract 2 electrons, electrons = = neutrons = =

10 10 CHEMISTRY ISOTOPES are Since isotopes have the same atomic number, however, since the atomic masses are different, Most elements exist as a mixture of several different. The molar mass of element is EXAMPLE VIII.2 Problem CALCULATING AVERAGE MOLAR MASSES Chlorine exists as a mixture of 75.77% Cl-35 and 24.23% Cl-37. If the precise molar mass of Cl-35 is g / mol and Cl-37 is g / mol, what is the average molar mass of the chlorine atoms? Solution mass of Cl-35 = ( ) x ( g / mol ) = g / mol mass of Cl-37 = ( ) x ( g / mol ) = g / mol total mass = g + g = g / mol If the exact masses of the isotopes are not given in the question, the atomic masses can be used instead. 35 Cl = 75.77% and 37 Cl = 24.23% average mass = ( ) + ( ) = g / mol

11 UNIT VIII ATOMS AND THE PERIODIC TABLE 11 C. The Electronic Structure of the Atom 1. When atoms are irradiated with energy, If the light emitted is passed through a prism and then onto photographic film, a is observed. In 1913, Niels Bohr proposed a model which explained the appearance of a hydrogen atom s line spectrum. He proposed that These energy states are Electrons could According to Bohr, the pattern of lines in the spectrum reflects the energy level pattern. The observed spectrum 2. The emerged from Bohr s theories of electron orbits. Several significant changes were made to the Bohr s basic ideas, the most notable being that

12 12 CHEMISTRY 11 An ORBITAL is 3. The is arranged as follows. Each dash represents the energy possessed by a particular orbital in the atom. The letters s, p, d, and f refer to four different types of orbitals. ENERGY LEVEL DIAGRAM FOR HYDROGEN A SHELL is For example, the 3 rd shell consists of the 3s, 3p, and 3d orbitals. A SUBSHELL is For example, the set of five 3d-orbitals in the 3 rd shell is a subshell. As can be seen on the above energy level diagram, (this is not true for atomes with more than one electron).

13 UNIT VIII ATOMS AND THE PERIODIC TABLE 13 The rules governing which types of orbitals can occur for a given energy level, and how many orbitals of a given type can exist, are: i) For a given value of n, n different types of orbitals are possible. for n = ; only the is possible. for n = ; the are possible. for n = ; the are possible. for n = ; the are possible. ii) An s-type subshell consists of s-orbital. A p-type subshell consists of p-orbitals. A d-type subshell consists of d-orbitals. An f-type subshell consists of f-orbitals. 4. The energy level diagram for hydrogen must be modified to describe any other atom. The modified diagram below

14 14 CHEMISTRY 11 ENERGY LEVEL DIAGRAM FOR POLYELECTRONIC ATOMS 5. An ELECTRON CONFIGURATION is a In particular, which orbitals in an atom contain electrons and how many electrons are in each orbital. The addition of electrons to the orbitals of an atom follows 2 simple rules. i) As the atomic number increases, electrons are added to the available orbitals. Electrons are The order in which orbitals are filled is

15 UNIT VIII ATOMS AND THE PERIODIC TABLE 15 ii) A maximum of electrons can be placed in each orbital. This means there can be a MAXIMUM of: electrons in an s-type subshell electrons in a p-type subshell electrons in a d-type subshell electrons in an f-type subshell. EXAMPLE VIII.3 Problem Solution WRITING ELECTRONIC CONFIGURATIONS Write the electronic configurations for He, Li, O, and Cl. He has electrons He = Li has electrons Li = O has electrons O = Cl has electrons Cl =

16 16 CHEMISTRY The following diagram Silicon, Si, has 14 electrons, First electrons fill the orbital and complete 1 st shell. Next electrons fill the and orbitals completing the 2 nd shell. Remaining electrons fill the orbital and partially fill the orbital. Si = Technetium, Tc, has 43 electrons, First electrons fill orbital. Next electrons fill and Next electrons fill and Next electrons fill,, and Remaining electrons fill and partially fill Tc =

17 UNIT VIII ATOMS AND THE PERIODIC TABLE The set of electrons belonging to a given atom can be divided into two subsets: the electrons and the electrons. The CORE of an atom is The OUTER electrons Since core electrons normally don t take part in chemical reactions, they are not always explicitly included when writing the electronic configuration of an atom. Core notation is In this notation, 8. There are two exceptions to the configurations of elements up to Kr. Instead of finding the actual configurations are Cr = [Ar] 4s 2 3d 4 and Cu = [Ar] 4s 2 3d 9

18 18 CHEMISTRY 11 Cr = and Cu = 9. There are two rules for writing the electronic configurations for ions. i) Negative Ions To write the electron configuration of a negative ion, O 2- = O 2- = O 2- = ii) Positive Ions Starting with the neutral configuration, If there are electrons in both the s- and p-orbitals of the outermost shell, Write core notation for the atom, remove electrons in the order: -electrons before -electrons before -electrons Sn 2+ Start with Sn = [Kr] 5s 2 4d 10 5p 2 remove 2 e - Sn 2+ =

19 UNIT VIII ATOMS AND THE PERIODIC TABLE 19 Sn 4+ Starting with [Kr] 5s 2 4d 10 5p 2 Sn 4+ = 10. Valence electrons are Valence electrons are all the electrons in the atom Al = [Ne] 3s 2 3p 1 Pb = [Xe] 6s 2 4f 14 5d 10 6p 2 Xe = [Kr] 5s 2 4d 10 5p 6 has valence electrons has valence electrons has valence electrons (noble gas configuration)

20 20 CHEMISTRY 11 D. Organizing the Elements The Periodic Table 1. Following Dalton s Atomic Theory, By 1817, chemists had discovered 52 elements and by 1863 that number had risen to In 1869 Russian chemist published a method of organizing the elements Mendeleev showed that He broke the list into a series of rows such that elements in one row were directly over elements with similar properties in other rows. He called each horizontal row a and each vertical column a. In certain cases, Mendeleev Mendeleev also left gaps in his table for elements which he believed were not yet discovered. He was so confident in his method of organization that When these elements were eventually discovered, they matched Mendeleev s predictions quite closely. At last with Mendeleev s Periodic Table,

21 UNIT VIII ATOMS AND THE PERIODIC TABLE As more and better data became available, chemists made a significant change to Mendeleev s method of organizing the elements. The modern periodic table is organized This solved problems The Periodic Law summarizes the organization of the periodic table. THE PERIODIC LAW 4. In the modern periodic table, a PERIOD is the A GROUP or FAMILY is the There are several special groups, rows, and blocks of elements. The are the main groups of elements. The are the central block of elements which separates the two blocks of the representative elements. The are the elements in the first column (except hydrogen).

22 22 CHEMISTRY 11 The are the elements in the second column. The are the elements in group 17 headed by fluorine. The are the elements in group 18 headed by helium. The and are the two rows below the main part of the table starting with lanthanum and actinium respectively. 5. Elements can also be classified according to their metallic character. The properties of metals

23 UNIT VIII ATOMS AND THE PERIODIC TABLE 23 The properties of nonmetals 6. There are some elements The nonmetals can be divided into two subgroups, A SEMICONDUCTOR is Semiconductors were formerly called or because they have properties which resemble metals more than nonmetals. The important difference is that the electrical conductivity of metals with increasing temperatures whereas the electrical conductivity of semiconductors with increasing temperature.

24 24 CHEMISTRY There are two important trends in the periodic table which exists among the elements: i) ii)

25 UNIT VIII ATOMS AND THE PERIODIC TABLE 25 E. Chemical Bonding 1. An ELECTROSTATIC FORCE is All chemical bonding is based on the following relationships of electrostatics: The greater the distance between two charged particles, The greater the charge on two particles, 2. Each period on the periodic table The 1 st shell has electrons and therefore the 1 st period has elements. The 2 nd shell has electrons and therefore the 2 nd period has elements. The 3 rd shell has electrons and therefore the 3 rd period has elements. (Note: for the purposes of this section, the transition metals, lanthanides, and actinides are IGNORED. Only the REPRESENTATIVE ELEMENTS will be considered.)

26 26 CHEMISTRY Going from left to right across a given period, the This increase in atomic number also brings an increase in the number of electrons surrounding the nucleus. All the electrons in a given shell can be assumed to have the same average distance from the nucleus. p+ 3p+

27 UNIT VIII ATOMS AND THE PERIODIC TABLE The shells surrounding the nucleus can be described as or An open shell is A closed shell is A closed shell has e.g. The 3 rd shell, Na to Ar, can hold a maximum of 8 electrons: 3s 2 3p 6. The atoms Na to Cl have less than 8 electrons in their 3 rd shell so they are OPEN. The atom Ar has its outermost shell full with 8 electrons therefore it is CLOSED. Previously VALENCE ELECTRONS were described as all the electrons in an atom excluding those in the core or filled d- or f-subshells. Valence electrons are. The NOBLE GASES have NO valence electrons and are NOT REACTIVE but F and Na HAVE valence electrons and ARE REACTIVE.

28 28 CHEMISTRY Isolated atoms have their electrons placed in s, p, d, and f orbitals; however, Only electrons are considered for bonding and the TRANSITION metals are ignored. There are a total of orbitals into which electrons can be placed (one s and three p orbitals). Each individual orbital holds up to electrons. Since electrons repel each other, Only after each orbital contains one electron will the The following or electron dot diagrams show how the valence electrons are distributed in an atom. The VALENCE (not valence electrons) of an atom = the number of electrons. Valence is sometimes called

29 UNIT VIII ATOMS AND THE PERIODIC TABLE In order to form a positive ion, an Li + energy Li + + e- IONIZATION ENERGY is (The electron is removed from the outermost shell and is always a valence electron unless the atom has a closed shell.) Ionization energy left to right across a period since The noble gas at the end of any period will Ionization energy top to bottom along a group since Ionization energy across a period and in a group.

30 30 CHEMISTRY 11 F. Types of Chemical Bonding 1. Atoms can form ions by either or electrons. Metal atoms generally form ions and nonmetal atoms form ions due to their difference in ELECTRONEGATIVITY. Electronegativity is Atoms of high electronegativity Electronegativity increases when the Electronegativity across a period and a group. In general, when an atom forms an ion, Na (1 valence e - ) e - + Na + (0 valence e - s like Ar) O (6 valence e - s) + 2e - O 2- (0 valence e - s like Ar) The most common charges found when going across the periodic table are shown below. The elements in group 14 (C, Si, Ge, Sn, and Pb) are not included because C, Si, and Ge do not form simple ionic compounds and Sn and Pb are metals which most readily form +2 ions and only rarely form +4 ions.

31 UNIT VIII ATOMS AND THE PERIODIC TABLE 31 Group Charge on ion 2. An IONIC BOND is formed It is formed when IONIC BONDS are formed when IONIC BONDS are very STRONG, so that compounds held together by ionic bonds have HIGH MELTING TEMPERATURES. In an IONIC SOLID there are Instead, there is a matrix of alternating positive and negative ions in three dimensions. Ionic solids are described as which are the lowest whole number ratio of positive to negative ions.

32 32 CHEMISTRY When an atom forms an ion, the resulting ion will be a different size than the corresponding neutral atom. When an atom gains electrons to form a ION, the does not change but since there are more electrons, the increases and the ion becomes than the neutral atom. When an atom loses electrons to form a ION, the between electrons decreases since there are fewer electrons. As such, and the ion becomes than the neutral atom. 4. In a COVALENT BOND Instead, the bond involves A covalent bond is formed when The states that atoms in groups 14 to 17 of the periodic table tend to form covalent bonds so as to have electrons in their valence shells.

33 UNIT VIII ATOMS AND THE PERIODIC TABLE 33 Atoms that form covalent bonds have relatively high electronegativities. They attract each other s electrons strongly but will not let go of their own electrons. This results in a tug of war and the electrons are shared in the bond. COVALENT BONDS are formed when a combines with a. 5. Oxygen atoms are 2 electrons short of a full shell. By sharing electrons and forming a, the atoms can form a full octet. Similarly, nitrogen atoms are 3 electrons short of a full shell. By sharing electrons and forming a, that atoms can form a full octet.

34 34 CHEMISTRY In covalent bonds between two different types of atoms the electrons This results in a bond where one end of the bond is slightly more negative (δ-) and the other end slightly more positive (δ+). Chemical bonds can be classified according to their difference in electronegativities. The following table lists the electronegativities of the elements H 1.0 Li 0.9 Na 0.8 K 0.8 Rb 0.7 Cs 0.7 Fr Be 1.2 Mg 1.0 Ca 1.0 Sr 0.9 Ba 0.9 Ra Sc 1.2 Y La-Lu Ac-No Zr Nb Mo Tc Ru Rh Pd Ag Cd 1.3 Hf 1.5 Ta 1.7 W 1.9 Re 2.2 Os 2.2 Ir 2.2 Pt 2.4 Au Hg B 1.5 Al C 1.8 Si Ti V Cr Mn Fe Co Ni Cu Zn Ga Ge 1.7 In 1.8 Tl Sn 1.8 Pb 3.0 N 2.1 P 2.0 As 1.9 Sb 1.9 Bi 3.5 O 2.5 S 2.4 Se 2.1 Te 2.0 Po 4.0 F 3.0 Cl 2.8 Br 2.5 I 2.2 At Type of Chemical Bond Electronegativity Difference Ionic Polar Covalent Covalent NaCl = = ( ) CH 4 = = ( ) F 2 = = ( )

35 UNIT VIII ATOMS AND THE PERIODIC TABLE The formula of covalently-bonded binary compounds can be predicted. Group Valence If we want to predict that formula of a compound between N and F. G. Writing Lewis Structures 1. LEWIS STRUCTURES (electron dot structures) are used to The symbol is used to denote the and dots are used indicate the To write the Lewis Structure for an atom, Li Ca B C N

36 36 CHEMISTRY The Lewis Structure of an ionic compound is written by: e.g. Draw the Lewis Structure for MgCl 2 3. Drawing Lewis Structures of covalent compounds that obey the octet rule follow a simple set of rules. Add one electron for each charge and subtract one electron for each charge. Determine which atoms are bonded together and Use the remaining valence electrons to Then (These non-bonding pairs of electrons are called.) If a central atom has less than an octet of electrons, Tidy up,

37 UNIT VIII ATOMS AND THE PERIODIC TABLE 37 Draw Lewis Structures for the following: NH 4 + CHO 2 - HOPO RESONANCE STRUCTURES exist when as in CHO - 2 and HOPO. 4. There are a number of atoms that violate the octet rule. In addition to H, the atoms are exceptions to the tendency for covalently-bonded atoms to complete their octet. These atoms have such a low electronegativity that they Be has 2 valence electrons and can only share 4 electrons (forming 2 bonds) while B and Al have 3 valence electrons and can only share 6 electrons (3 bonds).

38 38 CHEMISTRY 11 The Lewis Structure for BF 3 is A molecule in which one or more atoms (other than hydrogen) does not possess a full octet of electrons is called an molecule. 5. Elements in the 3 rd and 4 th periods of the periodic table Other than the fact that the central atom will end up with more than 8 electrons, the same rules are used to draw Lewis Structures. The Lewis Structure for PCl 5 is

39 UNIT VIII ATOMS AND THE PERIODIC TABLE 39 H. The Shape and Behaviour of Molecules 1. Lewis structures can be used to help visualize molecules in three dimension. Since all electrons carry the same charge, The valence electrons should be evenly spread out in regions of space around the central atom. This is the basis of the Summary of VSEPR shapes. Bonds Lone Pairs Shape Example Structure BeCl 2 BCl 3 CH 4 NH 3 H 2 O PCl 5 ClF 3

40 40 CHEMISTRY 11 SF 6 BrF 5 XeF 4 2. Polar bonds are a result of varying electronegativities among the elements. Since molecules usually possess more that one bond, If the polar bonds of a molecule are If the polar bonds are As a general rule, (Square planar molecules are an exception to this rule.)

41 UNIT VIII ATOMS AND THE PERIODIC TABLE Individual molecules are held together by covalent bonds between the atoms in the molecule. Such bonds are strong and are called. There are also weak forces that hold individual molecules next to other molecules. These are called There are two main types of van der Waals forces: Polar molecules are often referred to as because these molecules have a slightly positive and slightly negative end. As a result of the these dipoles, and they affect many of the properties of a compound such as boiling point There is a special case of dipole-dipole forces known as hydrogen-bonding. A HYDROGEN-BOND A hydrogen-bond is simply a particularly strong dipole-dipole force. H O H H O H O H H

42 42 CHEMISTRY 11 London Forces are the of van der Waals forces and are the result of. London forces are the weakest type of bonding force known. In general, London forces are always present, but are much weaker than covalent or ionic bonds. Hence, That is, London forces are important between the following closed-shell species: i) ii)

1. Following Dalton s Atomic Theory, 2. In 1869 Russian chemist published a method. of organizing the elements. Mendeleev showed that

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