Inorganic Chemistry Notes PDF

An atom is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number o...

An atom is the smallest unit of ordinary matter that forms a chemical element. Every solid, liquid, gas, and plasma is composed of neutral or ionized atoms. Every atom is composed of a nucleus and one or more electrons bound to the nucleus. The nucleus is made of one or more protons and a number of neutrons. Only the most common variety of hydrogen has no neutrons. More than 99.94% of an atom's mass is in the nucleus. The protons have a positive electric charge, the electrons have a negative electric charge, and the neutrons have no electric charge. If the number of protons and electrons are equal, then the atom is electrically neutral. If an atom has more or fewer electrons than protons, then it has an overall negative or positive charge, respectively – such atoms are called ions. The number of protons in the nucleus is the atomic number and it defines to which chemical element the atom belongs. For example, any atom that contains 29 protons is copper. The number of neutrons defines the isotope of the element. Atoms can attach to one or more other atoms by chemical bonds to form chemical compounds such as molecules or crystals. The ability of atoms to associate and dissociate is responsible for most of the physical changes observed in nature 1. Elements: are composed of extremely small particles called atoms. All atoms of a given element are identical. The atoms of one element are different from the atoms of all other elements. 2. Compounds are composed of atoms of more than one element. The relative number of atoms of each element in a given compound is always the same. Chemical reactions only involve the rearrangement of atoms. Atoms are not created or destroyed in chemical reactions. 2 Isotopes: Are two or more types of atoms that have the same atomic number (number of protons in their nuclei), and that differ in mass numbers due to different numbers of neutrons in their nuclei The electron configuration: Is the distribution of electrons of an atom or molecule (or other physical structure) in atomic or molecular orbitals Atomic number (Z) = number of protons in nucleus Mass number (A) = number of protons + number of neutrons = atomic number (Z) + number of neutrons Isotopes: are atoms of the same element (X) with different numbers of neutrons in the nucleus Mass Number A ZX Element Symbol Atomic Number 1 2 3 1H 1H (D) 1H (T) 235 238 92 U 92 U 2.3  Do You Understand Isotopes? How many protons, neutrons, and electrons are in 146 C? 6 protons, 8 (14 - 6) neutrons, 6 electrons How many protons, neutrons, and electrons are in 116 C? 6 protons, 5 (11 - 6) neutrons, 6 electrons 2.3 Atoms, Molecules and Ions A molecule is an aggregate of two or more atoms in a definite arrangement held together by chemical bonds H2 H2O NH3 CH4 A diatomic molecule contains only two atoms H2, N2, O2, Br2, HCl, CO A polyatomic molecule contains more than two atoms O3, H2O, NH3, CH4 A monatomic ion contains only one atom Na+, Cl-, Ca2+, O2-, Al3+, N3- A polyatomic ion contains more than one atom OH-, CN-, NH4+, NO3- Some Polyatomic Ions NH4+ ammonium SO42- sulfate CO32- carbonate SO32- sulfite HCO3- bicarbonate NO3- nitrate ClO3- chlorate NO2- nitrite Cr2O72- dichromate SCN- thiocyanate CrO42- chromate OH- hydroxide An ion is an atom, or group of atoms, that has a net positive or negative charge. cation – ion with a positive charge If a neutral atom loses one or more electrons it becomes a cation. 11 protons 11 protons Na 11 electrons Na + 10 electrons anion – ion with a negative charge If a neutral atom gains one or more electrons it becomes an anion. 17 protons 17 protons Cl 17 electrons Cl - 18 electrons A molecular formula shows the exact number of atoms of each element in the smallest unit of a substance An empirical formula shows the simplest whole-number ratio of the atoms in a substance molecular empirical H2O H2O C6H12O6 CH2O O3 O N2H4 NH2 Ionic compounds consist of a cation and an anion the formula is always the same as the empirical formula the sum of the charges on the cation and anion in each formula unit must equal zero The ionic compound NaCl Chemical Nomenclature Ionic Compounds often a metal + nonmetal anion (nonmetal), add “ide” to element name BaCl2 barium chloride K2O potassium oxide Mg(OH)2 magnesium hydroxide KNO3 potassium nitrate Formula of Ionic Compounds 2 x +3 = +6 3 x -2 = -6 Al2O3 Al3+ O2- 1 x +2 = +2 2 x -1 = -2 CaBr2 Ca2+ Br- 1 x +2 = +2 1 x -2 = -2 Na2CO3 Na+ CO32-  Transition metal ionic compounds indicate charge on metal with Roman numerals FeCl2 2 Cl- -2 so Fe is +2 iron(II) chloride FeCl3 3 Cl- -3 so Fe is +3 iron(III) chloride Cr2S3 3 S-2 -6 so Cr is +3 (6/2) chromium(III) sulfide  Molecular compounds nonmetals or nonmetals + metalloids common names H2O, NH3, CH4, C60 element further left in periodic table is 1st element closest to bottom of group is 1st if more than one compound can be formed from the same elements, use prefixes to indicate number of each kind of atom last element ends in ide  Molecular Compounds HI hydrogen iodide NF3 nitrogen trifluoride SO2 sulfur dioxide N2Cl4 dinitrogen tetrachloride NO2 nitrogen dioxide TOXIC! N2O dinitrogen monoxide Laughing Gas An acid can be defined as a substance that yields hydrogen ions (H+) when dissolved in water. HCl Pure substance, hydrogen chloride Dissolved in water (H+ Cl-), hydrochloric acid An oxoacid is an acid that contains hydrogen, oxygen, and another element. HNO3 nitric acid H2CO3 carbonic acid H2SO4 sulfuric acid A base can be defined as a substance that yields hydroxide ions (OH-) when dissolved in water. NaOH sodium hydroxide KOH potassium hydroxide Ba(OH)2 barium hydroxide  Orbitals Each orbital has room to hold a pair of electrons The final property of the electrons that we describe is the direction they are spinning, with their “north” pole pointing either up or down We show this by drawing the electrons as arrows pointing either up or down  Aufbau Principle: The Aufbau Principle (also called the building-up principle or the Aufbau rule) states that, in the ground state of an atom or ion, electrons fill atomic orbitals of the lowest available energy level before occupying higher-energy levels Pauli Exclusion Principle: no more than two electrons can occupy the same orbital and two electrons in the same orbital must have opposite spins Hund’s Rule: Electrons distribute themselves within a sub-level to maximize the number of unpaired electrons Writing orbital diagrams An orbital diagram shows every electron in an atom using an arrow: 1s 2s 2p Writing electron configurations A shorter way to display the arrangement of an atom’s electrons is its electron configuration: The configuration for the previous orbital diagram would be: 1s 2s 2p 1s2 2s2 2p4  There are four different kinds of orbitals, each having a unique shape These orbitals are designated by the letters s, p, d, and f. Which type of orbital an electron occupies depends on how much energy it has Each energy level corresponds to the electron occupying a particular orbital The higher the energy level, the larger the orbital Energy levels in detail The energy level an electron occupies is called its principal quantum number, designated by the letter n The lowest energy orbital an electron can occupy (n=1) is called its “ground state” Each energy level is divided into sublevels corresponding to the s, p, d, and f orbitals The number of sublevels in a principal energy level equals the quantum number n for that energy level The lowest energy level is n=1, which contains only 1 sublevel, called 1s The next energy level is n=2, which contains 2 sublevels, called 2s and 2p Each sublevel has room for different numbers of orbitals: s holds 1, p holds 3, d holds 5, f holds 7 Summary of energy levels and orbitals Principal energy level Sublevels Orbitals (n) 1 1s 1 2 2s, 2p 1+3=4 3 3s, 3p, 3d 1+3+5=9 4 4s, 4p, 4d, 4f 1+3+5+7=16 Metal Physical Properties: Lustrous (shiny) Good conductors of heat and electricity High melting point High density (heavy for their size) Malleable (can be hammered) Ductile (can be drawn into wires) Usually solid at room temperature (an exception is mercury) Opaque as a thin sheet (can't see through metals) Metals are sonorous or make a bell-like sound when struck Metal Chemical Properties: Have 1-3 electrons in the outer shell of each metal atom and lose electrons readily Corrode easily (e.g., damaged by oxidation such as tarnish or rust) Lose electrons easily Form oxides that are basic Have lower electronegativities Are good reducing agents Nonmetal Physical Properties: Not lustrous (dull appearance) Poor conductors of heat and electricity Nonductile solids Brittle solids May be solids, liquids or gases at room temperature Transparent as a thin sheet Nonmetals are not sonorous Nonmetal Chemical Properties: Usually have 4-8 electrons in their outer shell Readily gain or share valence electrons Form oxides that are acidic Have higher electronegativities Are good oxidizing agents Both metals and nonmetals take different forms (allotropes), which have different appearances and properties from each other. For example, graphite and diamond are two allotropes of the nonmetal carbon, while ferrite and austenite are two allotropes of iron. While nonmetals may have an allotrope that appears metallic, all of the allotropes of metals look like what we think of as a metal (lustrous, shiny). The Metalloids The distinction between metals and nonmetals is somewhat fuzzy. Elements with properties of both metals and nonmetals are called semimetals or metalloids. A stair-step line roughly divides metals from nonmetals on the periodic table. But, chemists recognize that naming one element a "metal" and the one next to it a "metalloid" is a judgement call. In truth, most metals display the properties of nonmetals under certain conditions, and nonmetals act like metals in some situations. Hydrogen is a good example of an element that acts as a nonmetal some times, but as a metal other times. Under normal conditions, hydrogen is a gas. As such, it acts like a nonmetal. But, under high pressure it becomes a solid metal. Even as a gas, hydrogen often forms the +1 cation (a metallic property). Yet, sometimes it forms the -1 anion (a nonmetal property). Groups 1 & 2, the Alkali Metals and the Alkaline Earth metals  Representative Elements Groups 1 and 2 Groups 1 and 2 are always found in nature combined with other elements. They’re called active metals because of their readiness to form new substances with other elements. They are all metals except hydrogen, the first element in Group 1. Although hydrogen is placed in Group 1, it shares properties with the elements in Group 1 and Group 17.  Alkali metals The Group 1 elements have a specific family name—alkali metals. All the alkali metals are silvery solids with low densities and low melting points. These elements increase in their reactivity, or tendency to combine with other substances, as you move from top to bottom.  Alkaline Earth Metals Next to the alkali metals are the alkaline earth metals. Each alkaline earth metal is denser and harder and has a higher melting point than the alkali metal in the same period. Alkaline earth metals are reactive, but not as reactive as the alkali metals. 1H Configuration 1s1 Can lose one electron to from stable ion H+ Can form single bond H2O, H2 2 He Configuration 1s2 Full Shell Inert Noble Gas Unlikely to form bonds or ions  3Li – Lithium – remember: can’t have s 3 (Pauli) Configuration 1s22s1 Easily ionized to Li+ (1s2) Alkali Metal Under standard conditions: lightest metal and least dense element Be- Beryllium 4 Configuration 1s22s2 Be2+ (1s2) stable ion Alkaline Earth Metal Toxic: replaces Magnesium from Magnesium activated enzymes due to stronger coordination ability 5B - Boron Number of Electrons: 5 Config 1s22s2 2p1 Forms stable covalently bonded molecular networks Mainly trivalent Lewis acidity of many of its compounds and multi-centre bonding Chemistry highly diverse and complex 6C - Carbon Config: 1s22s2 2p2 Diamond, graphite, amorphous, fullerenes (e.g., C60) Forms (usually) four bonds (tetravalent), see CH4 Non-metallic Carbon's unique characteristic of bonding to itself is responsible for complex molecules composed of long chains of carbon atoms, the skeleton of life 7N - Nitrogen Config 1s22s2 2p3 In compounds typically forms three bonds (trivalent), see NH3, N2 Non-metal N2: Gaseous, odourless, tasteless 78% of air is N2 Very inert at room temperature (and below) due to strong triple bond 8O Config 1s22s2 2p4 member of chalcogen group Normally considered divalent and highly electronegative O2: Colourless, odourless, highly reactive gas, Created biologically from CO2 by green photosynthesizing plants Paramagnetic (attracted by a magnetic field) 9F Config 1s22s2 2p5 Member of the halogen group Most reactive of all elements Forms F- readily highly electronegative! Does not exist in nature in the elemental state at all because of high reactivity 10 Ne - Neon Config 1s22s2 2p6 Full Outer shell Noble Gas Unlikely to form bonds or ions  Ionisation energy Ionization energy is the minimum amount of energy required to remove the most loosely bound electron of an isolated neutral gaseous atom or molecule. It is quantitatively expressed as X(g) + energy ⟶ X+(g) + e− Where, X is any atom or molecule, X+ is the ion with one electron removed, and e− is the removed electron  Atomic Radii a) Atomic radii generally decrease moving from left to right within the periods (nuclear charge keeps on increasing but electrons are added to the same shell), eg, going from Li (1s2 2s1) 157pm to F (1s2 2s22p6) 64pm; for both n = 2 b) Atomic radii generally increase down the group with increasing atomic number as electrons are occupying more and more distant electron shells, eg, going from Li (1s2 2s1) 157pm to Cs (1s2 2s22p63s23p64s23d104p65s24d105p66s2) 272pm c) There is a large increase as electrons go into next shell (like between He and Li or Ne to Na) d) All anions are larger than their parent atoms and all cations are smaller, compare Be2+(27 pm) and Be (112pm), I-(206pm) and I(133) – please note that ionic radius depends on coordination number of ion e) Ionic radii generally decrease with increasing positive charge on the same ion (Tl+, 164pm > Tl3+, 88pm) Atomic Radius: As you go down a group, the atomic radius increases with increasing energy level. Atomic radius decreases as you go left to right along a period because the greater nuclear charge pulls the electrons in closer to the nucleus. s-block p-block Electron Affinity: Is defined as the amount of energy released when an electron is attached to a neutral atom or molecule in the gaseous state to form a negative ion. X(g) + e− → X−(g) + energy. increases going left to right, and decreases going down. Electronegativity: The ability of an atom to attract electrons in a molecule. Electronegativity increases going left to right, and decreases going down.  Electronegativity The electronegativity of an atom is a measure of its power when in chemical combination to attract electrons to itself With few exceptions, electronegativity increases across the periodic table and decreases down a group, F is far more electronegative than I F is far more electronegative than Li A few definitions: van der Waals radius: The nonbonding radius of atoms Covalent radius: The radius of atoms covalently bonded to another atom; is smaller than the van der Waals radius Isoelectronic: Different ions that have the same number of electrons Ex: O2-, F-, Na+, Mg2+, Al3+ all have 10 electrons Ionization energy: The amount of energy required to remove an electron from an atom or ion. Electron affinity: The amount of energy released when an electron is added to an atom or ion. Periodic Properties of the Elements Electronegativity Ionization Energy & Electron Affinity Ionization Energy & Electron Affinity Electronegativity There are a number of atomic characteristics that either increase or decrease along the periodic table. Electron configurations of ions If electrons are added to make an anion, they fill the lowest energy levels first. (Auf bau principle) If electrons are removed to make a cation, they are taken from the highest energy levels first. Atom Ion Li 1s2 2s1 1s2 2s0 Li+ Fe [Ar] 4s2 3d6 [Ar] 4s0 3d6 Fe2+ Notice that the Fe3+ ion has the maximum [Ar] 4s0 3d5 Fe3+ multiplicity possible. Hence, its greater stability Fe3+: wrt Fe2+ 4s 3d Electron configuration exceptions and their ions Atom Ion Cr [Ar] 4s1 3d5 [Ar] 4s0 3d5 Cr+ [Ar] 4s0 3d3 Cr3+ [Ar] 4s0 3d0 Cr6+ Cu [Ar] 4s1 3d10 [Ar] 4s0 3d10 Cu+ [Ar] 4s0 3d9 Cu2+ Au [Xe] 6s1 4f14 5d10 [Xe] 6s0 4f14 5d10 Au+ [Xe] 6s0 4f14 5d8 Au3+ Why ionization energy decreases and atomic/ionic radius increases as go down a group: Shielding effect: The inner electron shells insulate the valence electrons from some of the electrical attraction with the positive charge of the nucleus. - + valence e- nucleus core electrons Why ionization energy increases and atomic/ionic radius decreases as go across a period: Increasing Effective Nuclear charge(Zeff): Electrons in the outermost energy levels do not effectively screen each other from an increasingly positive nucleus. - - - - nucleus - + + + + Zeff = Z - S + + + Li: Zeff = 3 – 2 = 1 N: Zeff = 7 – 2 = 5 The effective nuclear charge (often symbolized as Zeff or Z*) is the net positive charge experienced by an electron in a multi-electron atom. The term “effective” is used because the shielding effect of negatively charged electrons prevents higher orbital electrons from experiencing the full nuclear charge. The effective nuclear charge on an electron is given by the following equation: Zeff = Z – S where Z is the number of protons in the nucleus (atomic number), and S is the number of electrons between the nucleus and the electron (the number of nonvalence electrons).  Periodic Trends 1) Effective Nuclear Charge - Zeff 2s – 3s: Zeff greater for 3s probably due to actual higher overall charge  Shielding and effective nuclear charge Z* In polyelectronic atoms, each electron feels the attraction of the nucleus and the repulsion of the other electrons (both n and l must be taken into account) Each electron acts as a shield for electrons electrons farther away from the nucleus, reducing the attraction between the nucleus and the distant electrons Effective nuclear charge: Z* = Z – S (Z is the nuclear charge and S is the shielding constant) Example: Consider a neutral neon atom (Ne), a sodium cation (Na+), and a fluorine anion (F–). What is the effective nuclear charge for each? Solution: Start by figuring out the number of nonvalence electrons, which can be determined from the electron configuration. Ne has 10 electrons. The electron configuration is 1s22s2 2p6. The valence shell is shell 2 and contains 8 valence electrons. Thus the number of nonvalence electrons is 2 (10 total electrons – 8 valence). The atomic number for neon is 10, therefore: Zeff(Ne) = 10 – 2 = 8+ Flourine has 9 electrons but F– has gained an electron and thus has 10. The electron configuration is the same as for neon and the number of nonvalence electrons is 2. The atomic number for F– is 9, therefore: Zeff(F–) = 9 – 10 = 1 Sodium has 11 electrons but the Na+ ion has lost an electron and thus has 10. Once again, the electron configuration is the same as in the previous examples and the number of nonvalence electrons is 2 (by losing one electron, the valence shell becomes the n=3 shell). The atomic number for Na+ is 11, therefore: Zeff(Na+) = 11 – 10 = 1 In each of the above examples (Ne, F–, Na+) an atom has 10 electrons but the effective nuclear charge varies because each has a different atomic number. The sodium cation has the largest effective nuclear charge, which results in electrons being held the tightest, and therefore Na+ has the smallest atomic radius.  Dipole moment is a measure of the separation of positive and negative electrical charges within a system, that is, a measure of the system's overall polarity. μ = 𝛿.d Where: μ is the bond dipole moment, 𝛿 is the magnitude of the partial charges 𝛿+ and 𝛿–, And d is the distance between 𝛿+ and 𝛿–. molecules are affected by an electric field Which of the following molecules have a dipole moment? H2O, CO2, SO2, and CH4 O S O H H O H H C H O C O H  Polarity * Polarity is a separation of electric charge leading to a molecule or its chemical groups having an electric dipole moment, with a negatively charged end and a positively charged end. * Polar molecules must contain polar bonds due to a difference in electronegativity between the bonded atoms. A polar molecule with two or more polar bonds must have a geometry which is asymmetric in at least one direction, so that the bond dipoles do not cancel each other. * Polar molecules interact through dipole–dipole intermolecular forces and hydrogen bonds. Polarity underlies a number of physical properties including surface tension, solubility, and melting and boiling points.  Chemical forces An intramolecular force (or primary forces) is any force that binds together the atoms making up a molecule or compound, not to be confused with intermolecular forces, which are the forces present between molecules  Ionic bonding Ionic bonding is a type of chemical bonding that involves the electrostatic attraction between oppositely charged ions, or between two atoms with sharply different electronegativities, and is the primary interaction occurring in ionic compounds. It is one of the main types of bonding along with covalent bonding and metallic bonding  Covalent Bond A covalent bond is a chemical bond that involves the sharing of electron pairs between atoms. These electron pairs are known as shared pairs or bonding pairs, and the stable balance of attractive and repulsive forces between atoms, when they share electrons, is known as covalent bonding.  Hydrogen bonding A hydrogen bond is the attraction between the ione pair of an electronegative atom and a hydrogen atom that is bonded to an electronegative atom, usually nitrogen, oxygen, or fluorine. The hydrogen bond is often described as a strong electrostatic dipole–dipole interaction  Ion–dipole and ion–induced dipole forces Ion–dipole and ion–induced dipole forces are similar to dipole–dipole and dipole–induced dipole interactions but involve ions, instead of only polar and non-polar molecules. Ion–dipole and ion–induced dipole forces are stronger than dipole–dipole interactions because the charge of any ion is much greater than the charge of a dipole moment. Ion–dipole bonding is stronger than hydrogen bonding. An ion–dipole force consists of an ion and a polar molecule interacting.  Van der Waals forces Is a distance-dependent interaction between atoms or molecules. Unlike ionic or covalent bonds, these attractions do not result from a chemical electronic bond; they are comparatively weak and therefore more susceptible to disturbance. The Van der Waals force quickly vanishes at longer distances between interacting molecules.  Resonance In chemistry, resonance is a way of describing bonding in certain molecules or ions by the combination of several contributing structures (or forms, also variously known as resonance structures or canonical structures) into a resonance hybrid (or hybrid structure) in valence bond theory. It has particular value for describing delocalized electrons within certain molecules or polyatomic ions where the bonding cannot be expressed by one single Lewis structure. Resonance energy Is the energy released by the conjugated system due to delocalisation of electron  Melting point The melting point (or, rarely, liquefaction point) of a substance is the temperature at which it changes state from solid to liquid. At the melting point the solid and liquid phase exist in equilibrium. The melting point of a substance depends on pressure and is usually specified at a standard pressure such as 1 atmosphere or 100 kPa.  Boiling point The boiling point of a substance is the temperature at which the vapor pressure of a liquid equals the pressure surrounding the liquid and the liquid changes into a vapor. The boiling point of a liquid varies depending upon the surrounding environmental pressure. A liquid in a partial vacuum has a lower boiling point than when that liquid is at atmospheric pressure. A liquid at high pressure has a higher boiling point than when that liquid is at atmospheric pressure. For example, water boils at 100 °C (212 °F) at sea level, but at 93.4 °C (200.1 °F) at 1,905 meters (6,250 ft) altitude. For a given pressure, different liquids will boil at different temperatures.  Solubility Solubility is the property of a solid, liquid or gaseous chemical substance called solute to dissolve in a solid, liquid or gaseous solvent. The solubility of a substance fundamentally depends on the physical and chemical properties of the solute and solvent as well as on temperature, pressure and presence of other chemicals (including changes to the pH) of the solution.  Paramagnetism and Diamagnetism Paramagnetism is a form of magnetism whereby some materials are weakly attracted by an externally applied magnetic field, and form internal, induced magnetic fields in the direction of the applied magnetic field. In contrast with this behavior, diamagnetic materials are repelled by magnetic fields and form induced magnetic fields in the direction opposite to that of the applied magnetic field If all electrons in the particle are paired, then the substance made of this particle is diamagnetic; if it has unpaired electrons, then the substance is paramagnetic.  Orbital hybridisation Orbital hybridisation is the concept of mixing atomic orbitals into new hybrid orbitals (with different energies, shapes, etc., than the component atomic orbitals) suitable for the pairing of electrons to form chemical bonds in valence bond theory. For example, in a carbon atom which forms four single bonds the valence-shell s orbital combines with three valence-shell p orbitals to form four equivalent sp3 mixtures which are arranged in a tetrahedral arrangement around the carbon to bond to four different atoms. Hybrid orbitals are useful in the explanation of molecular geometry and atomic bonding properties and are symmetrically disposed in space. Usually hybrid orbitals are formed by mixing atomic orbitals of comparable energies  Types of hybridisation 1. sp3 For a tetrahedral coordinated carbon (e.g., methane CH4), the carbon should have 4 orbitals with the correct symmetry to bond to the 4 hydrogen atoms. Carbon's ground state configuration is 1s2 2s2 2p2 2. sp2 For example, ethene (C2H4) has a double bond between the carbons. For this molecule, carbon sp2 hybridises, because one π (pi) bond is required for the double bond between the carbons and only three σ bonds are formed per carbon atom. In sp2 hybridisation the 2s orbital is mixed with only two of the three available 2p orbitals, usually denoted 2px and 2py. The third 2p orbital (2pz) remains unhybridised. 3. sp The chemical bonding in compounds such as alkynes with triple bonds is explained by sp hybridization. In this model, the 2s orbital is mixed with only one of the three p orbitals, Coordination numb Shape Hybridisation Examples er 2 Linear sp hybridisation CO2 (180°) 3 Trigonal planar sp2 hybridisation BCl3 (120°) 4 Tetrahedral sp3 hybridisation CCl4 (109.5°) 4 Square planar sp2d hybridisation PtCl42− Trigonal bipyramidal Fe(CO)5 5 sp d hybridisation 3 Square pyramidal MnCl52− 6 Octahedral sp3d2 hybridisation Mo(CO)6 Shapes of simple covalent molecules