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CHM 1301
GENERAL CHEMISTRY I
FIRST PART LECTURE NOTES
GENERAL AND INORGANIC CHEMISTRY
Atom An atom is known as the smallest particle of an element which can take part in a chemical reaction. It comprises of tiny particles called electrons, protons and neutrons. These three particles are present in all atoms except for hydrogen with only one proton and one electron. Molecule This is known as the smallest particles of a substance, either an element or a compound which is capable of independent existence. Atomicity is known as the number of atoms in a molecules of a given element. Example, the atomicity of O 2 and H 2 is two and is known as diatomic. Others such as phosphorus and sulphur exist as polyatomic molecules. The molecules of helium and neon are monoatomic because they exist independently as a single atom. Mole DEFINITION: Mole is the amount of substance in grams that has the same number of particles as there are atoms in 12 grams of carbon-12. One mole of ethanol, for example, contains the same number of ethanol molecules as there are carbon atoms in 12 g of carbon-12. The number of atoms in a 12-g sample of carbon-12 is called Avogadro’s number (to which we give the symbol NA ). Recent measurements of this number gives the value 6.02x10^23. A mole of a substance contains Avogadro’s number (6.02x10^23 ) of molecules. The molar mass of a substance is the mass of one mole of the substance. Carbon- has a molar mass of exactly 12 g/mol, by definition. Problems (a) What is the number of moles in 35.0g of CuSO4? Moles = Mass M.M = 35.0/ (63.5 + 32 +16x4) = 0.219 mol [Cu = 63.5, S = 32, O = 16 ] (b) What is the number of moles in 75.0mg of CaSO 4 .2H 2 O? [Ca = 40, S = 32, O = 16, H = 1] Moles = Mass M.M
= 0.075/ (40 + 32 +16 x4 + 18x2) = 4.36x10-4^ mol How many molecules of ethanol are there in a 0.500 dm3 of ethanol (CH3CH2OH) liquid? The density of ethanol is 0.789 g cm- Number of molecules = moles x 6.022 x 10^23 = 8.576 x 6.022 x 10^23 = 5.16 x10^24 (to 3 sig fig) Mass = density x Volume = 0.789 g cm-3^ x 500 cm^3
= 394.5g ethanol Moles = Mass M.M = 394.5/ 46. = 8.576 mol Calculation of percentage composition by mass Example: Calculate the percentage composition by mass of each element present in Na 2 SO 4. Molar mass of Na 2 SO 4 = 2 x 23+1 x 32+ 4 x 16 = 46 + 32 + 64 = 142 gmol-^1 [Na = 23, O = 16, S = 32] Solution: Sodium = 14246 x 100% = 32.39%
Sulfur = 14232 x 100% = 22.5%
Oxygen = 14232 x 100% = 45 .07%
Atomic Theory in the Nineteenth Century Introduction: The earliest recorded discussion of the basic structure of matter comes from ancient Greek philosophers, the scientists of their day. In the fifth century BC, Leucippus and Democritus argued that all matter was composed of small, finite particles that they called atomos , a term derived from the Greek word for “indivisible.” They thought of atoms as moving particles that differed in shape and size, and which could join together. Later, Aristotle and others came to the conclusion that matter consisted of various combinations of the four “elements”—fire, earth, air, and water—and could be infinitely divided. Interestingly, these philosophers thought about atoms and “elements” as philosophical concepts, but apparently never considered performing experiments to test their ideas. The Aristotelian view of the composition of matter held sway for over two thousand years, until English school teacher John Dalton helped to revolutionize chemistry with his hypothesis that the behavior of matter could be explained using an atomic theory. First published in 1807, many of Dalton’s hypotheses about the microscopic features of matter are still valid in modern atomic theory. Here are the postulates of Dalton’s atomic theory.
- Matter is composed of exceedingly small particles called atoms. An atom is the smallest unit of an element that can participate in a chemical change.
- An element consists of only one type of atom, which has a mass that is characteristic of the element and is the same for all atoms of that element
- Atoms of one element differ in properties from atoms of all other elements.
- A compound consists of atoms of two or more elements combined in a small, whole- number ratio. In a given compound, the numbers of atoms of each of its elements are always present in the same ratio
- Atoms are neither created nor destroyed during a chemical change, but are instead rearranged to yield substances that are different from those present before the change.
- 1.602x10-19C. He calculated the mass of the electron as follows;
Mc = (^) (charge/mass)charge = - 1.602x10-19C
1.759x10^8 Cg-1^ = 9.107x10-28g Scientists had now established that the atom was not indivisible as Dalton had believed, and due to the work of Thomson, Millikan, and others, the charge and mass of the negative, subatomic particles - the electrons - were known. However, the positively charged part of an atom was not yet well understood. In 1904, Thomson proposed the “plum pudding” model of atoms, which described a positively charged mass with an equal amount of negative charge in the form of electrons embedded in it, since all atoms are electrically neutral (Fig. 1).
Figure 1: Thomson’s plum pudding description of the atom The next major development in understanding the atom came from Ernest Rutherford , a physicist from New Zealand who largely spent his scientific career in Canada and England. He performed a series of experiments using a beam of high-speed, positively charged alpha particles (α particles) that were produced by the radioactive decay of Radium; α particles consist of two protons and two neutrons Rutherford and his colleagues Hans Geiger (later famous for the Geiger counter) and Ernest Marsden aimed a beam of α particles, the source of which was embedded in a lead block to absorb most of the radiation, at a very thin piece of gold foil and examined the resultant scattering of the α particles using a luminescent screen that glowed briefly when hit by an α particle. What did they discover? Most particles passed right through the foil without being deflected at all. However, some were diverted slightly, and a very small number were deflected almost straight back toward the source (Figure 2).
Figure 2: Rutherford alpha paricle scttering experiment Here is what Rutherford deduced: Because most of the fast-moving α particles passed through the gold atoms undeflected, they must have traveled through essentially empty space inside the atom. Alpha particles are positively charged, so deflections arose when they encountered another positive charge (like charges repel each other). Since like charges repel one another, the few positively charged α particles that changed paths abruptly must have hit, or closely approached, another body that also had a highly concentrated, positive charge. Since the deflections occurred a small fraction of the time, this charge only occupied a small amount of the space in the gold foil. Analyzing a series of such experiments in detail, Rutherford drew two conclusions:
- The volume occupied by an atom must consist of a large amount of empty space.
- A small, relatively heavy, positively charged body, the nucleus , must be at the center of each atom. Conclusions: Rutherford concluded that: Positive charges (i.e protons) are concentrated in a dense central core within the atom called the nucleus. Rutherford model of the atom: According to the Rutherford, an atom consists of a central nucleus surrounded by revolving electrons
Bohr’s Model of The Atom In 1913, Neil Bohr proposed a model for the H-atom based on thje following postulates: i. The electron in the H-atom moves in the circular orbits round the nucleus. ii. Each orbit has a fixed energy value iii. As long as an electron stays in a given orbit, it does not radiate energy. iv. Radiation is emitted when an electron falls froms higher to a lower orbit and absorbed wheen an electron jumps from a lower to a higher level v. The energy of radiation emitted or absorbed is equal to to the energy difference between the two levels
- Precipitation reactions: In this type two soutions of ionic substances are mixed and a solid of ionic substance is precipitated.
NaCl(aq) + AgNO3(aq) NaNO3(aq) + AgCl(s)
- Aid-base reaction: This type of reaction involve transfer of proton between the reactants;
HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l)
- Oxidation- reduction reaction: This involve the transfer of electron(s) between the reactants.
Fe(s) CuSO4(aq) FeSO4(aq) + Cu(s)
Writing and balancing chemical equation:
- The substances undergoing reaction are called reactants , and their formulas are placed on the left side of the equation.
- The substances generated by the reaction are called products , and their formulas are placed on the right sight of the equation.
- Plus signs (+) separate individual reactant and product formulas, and an arrow (⟶) separates the reactant and product (left and right) sides of the equation.
- The relative numbers of reactant and product species are represented by coefficients (numbers placed immediately to the left of each formula). A coefficient of 1 is typically omitted. Examples: (a) Write a balanced chemical equation for the combustion of methane.
CH4(g) + 2O2(g) CO2(g) + H 2 O(l)
(b) Write a balanced equation for the reaction of molecular nitrogen (N 2 ) and oxygen (O 2 ) to form dinitrogen pentoxide. Solutio
2N2(g) + 5O2(g) 2N 2 O5(g)
Atomic Number, Mass Number and Isotopes Atomic Number (Z) is the number of proton in the nucleus of an atom of an element. In a neutral aton the number of proton is equal to the number of electrons. So atomic number also indicate the number of electrons in the atom. Mass number (A) is the total number of neutrons and protons present in the nucleus of an atom of an element. Number of neutrons = mass number – atomic number. N = A – Z Isotopes are atoms that have the same atomic number but different mass numbers. For examople hydrogen has three isotopes as given in the table below:
Isotope Name Number of Neutrons
(^11) H
Protium 0
12 H^ Deuterium^1 (^31) H^ Tritium^2
Quantum Numbers These are numbers used to completely identify an electron in an atom. (a) Principal quantum number (n) It has integral values of n = 1,2,3,...∞ which indicate the nergy levels of the electron. Example for n = 3 implies that, the electron is in the third energy level. N = ∞ indicates that the electron has beeen completely removed from the atom. (b) Subsidiary or azimuthal or angular momentum quantum number (l) It has integral values ranging from 0 n-1, which indicate the type of orbital containing an electron as given in the table below: l value Orbital 0 s 1 p 2 d 3 f Now let us consider the types of orbitals in various energy levels. For the first energy level, n = 1. The only possible value of l is 1-1 = 0. So, 0nly an s-orbital is present in the first energy level. For the second energy level, n = 2. The only possible values of l are 0,... (n-1) = 0,1. so, there are s and p-orbitals in the second energy level. Energy level (n) Possible values of l Orbitals in the energy level 1 0 s 2 0,1 s,p 3 0,1,2 s,p,d 4 0,1,2,3 s,p,d,f
All energy levels from n = 4,... ∞ have s,p,d and f orbitals. Magnetic quntum number (ml) For any value of l, ml has the integral values – l,…..+l which gives the number of orbital types in energy level. For example, in the third energy level, l = 0,1,2. For l = 0, ml = 0. There is only s-orbital. l =1, ml = -1,0,+1. There are three p-orbitals
Rules Governing the Filling of Electrons in Orbital The electron configuration of an element is the arrangement of electrons in its atomic orbitals when it is at the ground state.
The Aufbau principle Aufbau is a German word, which means “build-up” The Aufbau principle states that electrons at their ground state are added to an atom, one at a time, starting with the lowest energy orbital until all orbitals are filled. The Aufbau diagram is given below.
Figure 5: Aufbau diagram Pauli Exclusion Principle The Pauli Exclusion Principle states that no two electrons in the same atom can have the same values for all the four quantum numbers. Hund’s rule Hund’s rule states that electrons occupy degenerate orbitals (i.e. orbitals with equal energy) singly before each levels singly while maintaining parallel spins before they pair up. The table below shows the maximum number of electrons an orbital can accomodate Orbital Number of electrons accomadated s 1 or 2 p 1 to 6 d 1 to 10 f 1 to 14
Problems: Write the ground state electron configuration of the following atoms with the indicated atomic numbers and determine the groups, periods and blocks they belong to:
(a) 2 He (b) 5 N (c) 10 Ne (d) 12 Mg
Periodicty or Periodic Properties A periodic property is any property which varies systematically from one element to the next across the period and down the group in the periodic table.
Generally, a periodic property which increases along the period decreases down the group and vice-versa. Some of the periodic properties include: Atomic radius, ionization energy, electron affinity, electronegativity, density, melting point and boiling point.
a) Atomic radius is defined as the closest approach to another identical atom. When an atom is in a bonding situation, the atomic radius is the same as the covalent radius. The covalent radius is defined as half the distace between two covalently bonded atoms. When an atom is not in bonding situation, the atomic radius is the same as the Vander Waal’s radius, which is defined as the distance between two identical atoms which are not chemically bonded.
Figure 6: Atomic radius Variation of atomic radius across the period and down the group As we move from left to right across the period, the atoms become smaller due to increase effective nuclear charge and thus, atomic radius decrease across the period. While on moving down the groups the atomic radius increase because of the decrease effective nuclear charge due to the increase in the number of shells. b) Ionization energy is the minimum energy required to remove an electron completely from the atom. Variation of ionization energy across the period and down the group Moving from left to right across a period , effective nuclear charge inreases, this causes the outermost electrons are more tightly held making them harder to be removed. The ionization energy therefore, increase across the period. While as we move down the group, the atoms become bigger. The outer electrons are held less tightly at distances further away from the nucleus and thus, the ionization energy decrease down the group. c) Electron affinity is the energy change when an electron is added to an isolated gaseous atom to form a negatively charged ion. Variation of electron affinity across the period and down the group
The sp^3 hybrid orbitals are equivalent and degenerate. They differ only in their orientation in space.
2. sp^2 hybridization: sp^2 hybrid orbitals formed from the combination of one s-orbital and two p-orbitals. Example Boron has 3 unpaired electrons available for bonding once a 2s-electron is promoted to an empty 2p-orbital.The s-orbital and two of the p-orbitals hybridize to form sp^2 orbitals. The three sp^2 orbitals lie in a 120o^ to minimize electron repulsion. 3. sp hybridization: sp hybrid orbitals form from the combination of one s-orbital and 1 p-orbital. Carbon can also form sp hybrid orbitals. Acetylene (C 2 H 2 ) with C-C triple bond. 4. sp^3 d Hybridization: In this type of hybridization, one ‘s’, three ‘p’ and one ‘d’ of the same principal quantum number mix to give five sp^3 d hybridized orbitals.
Example: In PCl 5 , P atom is undergoing sp^3 d hybridization. The shape of the molecule is trigonal bipyramidal.
The types of hybridization and shapes of molecules Hybridization Number of atoms linked to central atom
Shape of molecule or ion
Example
Sp 2 Linear HgCl 2 , BeCl 2 , C 2 H2, CO 2 sp^2 3 Trigonal planar BF3, CO 32 - , SO 3 sp^2 2 Angular SnCl 2 , NO 2 - sp^3 4 Tetrahedral CH 4 , BF 4 - ,ClO 4 - sp^3 3 Trigonal pyramidal NH 3 , PCl 3 , ClO 3 - sp^3 2 Angular H 2 O, ClO 2 - , H 2 S
sp^3 d 5 Trigonal bipyramidal
PCl 5 , PF 5
sp^3 d 4 Distorted tetrahedral
SF 4 , IF 4 +
sp^3 d 3 T-Shaped ClF 3 , BrF 3 sp^3 d 2 Linear XeF 2 , I 3 - sp^3 d^2 6 Octahedral SF 6 , PF 6 - sp^3 d^2 5 Square pyramidal IF 5 sp^3 d^2 4 Square planar XeF 4 , ICl 4 -
Example: Give the hybridization of atomic orbital of nitrogen in (a) NO 2 +^ (b) NO 3 -^ (c) NH 4 + Solution: (a) NO 2 +^ Number of valence electron of nitrogen = 5 n = 5-1 = 4 (due to one +ve charge) n 2 =^
4 2 = 2, thus,^ the hybridization = sp (b) In NO 3 -^ n = 5+1 = 6 (due to one – ve charge) n 2 =^
6 2 = 3, thus the hybridization = sp
2
(c) In NH 4 +^ n = 5+4-1 (due to four hydrogen atoms and one positive charge) n 2 =^
8 2 = 4, thus the hybridization = sp
3
Example: find the molecular shapes of the following compounds (i) SF 4 (ii) CF 4 (iii) XeF 4 (i) In SF 4 n = 6+4 = 10 n 2 =^
10 2 = 5, thus the hybridization = sp
(^3) d, hence the shape is trigonal bipyramidal
(ii) CF 4 n = 4+4 = 8 n 2 =^
8 2 = 4, thus the hybridization of^ C = sp
(^3) , hence the shape is tetrahedral
(iii) XeF 4 n = 8+4 = 12 n 2 =^
12 2 = 6, thus the hybridization = sp
(^3) d (^2) , hence the shape is square planar
CHEMICAL EQUATIONS AND STOICHIOMETRY A balanced chemical equation provides a great deal of information in a very succinct format. Chemical formulas provide the identities of the reactants and products involved in the chemical change, allowing classification of the reaction. Coefficients provide the relative numbers of these chemical species, allowing a quantitative assessment of the relationships between the amounts of substances consumed and produced by the reaction. These quantitative relationships are known as the reaction’s stoichiometry , a term derived from the Greek words stoicheion (meaning “element”) and metron (meaning “measure”). In this topic, the use of balanced chemical equations for various stoichiometric applications is explored.
Mass = moles x M.M = 0.0328 x 44. =1.44g
- What mass of magnesium oxide would be produced from 16 g of magnesium in the reaction between magnesium and oxygen?
a. Write the full balanced equation
2Mg(s) + O2(g) → 2MgO(s)
b. Read the equation in terms of moles
2 moles of magnesium reacts to give 2 moles of magnesium oxide
iii Convert the moles to masses using the RMM values
∴ (2 x 24g) of magnesium gives 2 x (24+16)
= 80 g of Magnesium oxide ∴ 16 g of magnesium gives = 26.7 g of Magnesium oxide
- What volume of oxygen would react with 16 g of magnesium in the above reaction?
In this case the oxygen is a gas so the volume of each mole is 22,400 cm^3 at room temperature and pressure and you do not have to worry about the molecular mass of the gas.
From the equation:
2 moles of Mg reacts with 1 mole of O 2
∴2 x 24 g of Mg reacts with 1 x 22400 cm^3 of O2(g)
∴ 16 g of Mg reacts with 1×22400cm3 ×16g 2x24g = 7466.7cm^3 of oxygen
- What mass of lead (II) sulphate would be produced by the action of excess dilute sulphuric acid on 10 g of lead nitrate dissolved in water? Pb(NO 3 )2(aq) + H 2 SO4(aq) → PbSO4(s) + 2HNO3(aq) ∴ 1 mole of lead nitrate gives 1 mole of lead sulphate ∴ 331 g of lead nitrate gives 303 g of lead sulphate
∴ 10 g of lead nitrate gives 303gx10g of leadsulphate 331 g = 9.15 g of lead sulphate
- What is the total volume of gas produced by the action of heat on 1 g of silver nitrate? 2AgNO3(s) → 2Ag(s) + 2NO2(g) + O2(g) 2 moles of silver nitrate give 2 moles of nitrogen dioxide gas plus 1 mole of oxygen gas = 3 moles of gas ∴ 2 x 170 g of silver nitrate give 3 x 24000 cm3 of gas
∴1 g of silver nitrate gives 3×22400cm3 ×1g of gas2 x170g = 197.65 cm^3 of gas
CHEMICAL BONDING AND INTERMOLECULAR FORCES When two atoms of the same or different elements approach each other, the energy of the combination of the atoms becomes less than the sum of the energies of the two separate atoms at a large distance. We say that the two atoms have combined or a bond is formed between the two. The bond is called a chemical bond. Thus a chemical bond may be visualised as an effect that leads to the decrease in the energy. The combination of atoms leads to the formation of a molecule that has distinct properties different from that of the constituent atoms.
A question arises, “How do atoms achieve the decrease in energy to form the bond”. The answer lies in the electronic configuration. As you are aware, the noble gases do not react with other elements to form compounds. This is due to their stable electronic configuration with eight electrons (two in case of helium) in their outermost shells. The formation of a bond between two atoms may be visualised in terms of their acquiring stable electronic configurations. That is when two atoms (other than that of noble gases) combine they will do so in such a way that they attain an electronic configuration of the nearest noble gas.
The stable electronic configuration of the noble gases can be achieved in a number of ways; by losing, gaining or sharing of electrons. Accordingly, there are different types of chemical bonds, like,
Ionic or electrovalent bond
Covalent bond
Co-ordinate covalent bond
In addition to these we have a special kind of bond called hydrogen bond. Let us discuss about different types of bonds, their formation and the properties of the compounds so formed.
IONIC OR ELECTROVALENT BOND
- They are generally soluble in water and less soluble in non-polar solvents like ether, alcohol, etc.
- They conduct electricity when in molten state or in aqueous solutions. Kossel’s theory explains bonding quite well but only for a small class of solids composed of electropositive elements of Group 1 and 2 with highly electronegative elements. Secondly, this theory is incapable of explaining the formation of compounds like, SO 2 or O 2 , etc. For example in case of O 2 , there is no reason to expect that one atom of oxygen would lose two electrons while the other accepts them. The problem was solved by Lewis theory of covalent bonding.
COVALENT BONDING Like Kossel, Lewis also assumed that atoms attain noble gas electronic configuration in the process of bond formation. However, the way the noble gas electronic configuration is achieved, is different. Lewis proposed that this is achieved by “sharing of a pair of electrons” between the two atoms. Both the atoms contribute an electron each to this pair. For example, two hydrogen atoms form a molecule by sharing a pair of electrons. If electrons are indicated as dots, formation of hydrogen molecule can be shown as
H. +. H H : H H—H
This shared pair of electrons contributes towards the stability of both the atoms and is said to be responsible for ‘bonding’ between the two atoms. Such a bond is called covalent bond and the compounds so obtained are called covalent compounds. In the process of suggesting the process of chemical bonding Lewis provided a very convenient way of representing bonding in simple molecules. This is called Lewis electron-dot structures or simply Lewis structures.
In Lewis structure each element is represented by a Lewis symbol. This symbol consists of the normal chemical symbol of the element surrounded by number of dots representing the electrons in the valence shell. Since the electrons are represented by dots, these are called electron-dot structures. The Lewis symbols of some elements are as:
You may note here that while writing the Lewis symbols, single dots are placed first on each side of the chemical symbol then they are paired up. The Lewis structure of a molecule is written in terms of these symbols. Example the Lewis dot structure for the formation of some substances are as follows:
Polar Covalent Bond
In a chemical bond the shared electron pair is attracted by the nuclei of both the atoms. When we write the electron dot formula for a given molecule this shared electron pair is generally shown in the middle of the two atoms indicating that the two atoms attract it equally. However, actually different kinds of atoms exert different degrees of attraction on the shared pair of electrons. A more electronegative atom has greater attraction for the shared pair of electrons in a molecule. As a consequence in most cases the sharing is not equal and the shared electron pair lies more towards the atom with a higher electronegativity. For example, in HCl, the shared pair of electron is attracted more towards more electronegative chlorine atom. As a result of this unequal sharing of the electron pair , the bond acquires polarity or partial ionic character.
Coordinate Covalent Bond You have learnt that in the formation of a covalent bond between the atoms, each atom contributes one electron to the shared electron pair, However, in some cases both the
