Chemical Bonding PDF

Chemical Bonding: Chemical Formulas **Definition: A compound is a substance that is made up of two or more different elements combined together chemically.** When elements combine to form compounds, there are attractive forces that hold the atoms together in the new substance. These forces are calle...

Chemical Bonding: Chemical Formulas **Definition: A compound is a substance that is made up of two or more different elements combined together chemically.** When elements combine to form compounds, there are attractive forces that hold the atoms together in the new substance. These forces are called chemical bonds. The Octet Rule **Definition: Octet Rule – When bonding occurs, atoms tend to reach an electron arrangement with eight electrons in the outermost energy level.** The noble gases (helium, neon, argon) are unreactive compounds (inert). All of the noble gases consist of single atoms. These elements do not react because they have stable electron configurations. All noble gases, except helium, have eight electrons in their outermost energy level. When elements react together, they try to rearrange their configurations to have 8 electrons in the outermost energy level. Exceptions to the Octet Rule - Transition Metals usually do not obey the octet rule. - Hydrogen, Lithium and Beryllium tend to achieve two electrons in the outermost energy level instead of eight. Ionic Bonding – Transfer of Electrons **Definition: An ion is a charged atom or group of atoms.** Atoms can obtain the electron configuration of a noble gas by gaining or losing an electron. - The elements in group I of the periodic table tend to lose one electron to form an ion with one positive charge. - The elements in group II of the periodic table tend to lose two electrons to form an ion with two positive charges. - In a positive ion (cation), since the atom has lost electrons, there are more protons in the nucleus than there are electrons in orbit around it. - The elements in Group VI of the Periodic Table lose two electrons to form an ion with two negative charges. - The elements in Group VII of the Periodic Table lose one electron to form an ion with one negative charge. A negative ion is called an anion. **Definition: An ionic bond is the force of attraction between oppositely charges ions in a compound. Ionic bonds are always formed by the complete transfer of electrons from one atom to another.** 1 How to show the Formation of Ionic Bonding Dot and Cross Diagram - The outer electrons of one atom are represented by dots while the electrons on the other atom are denoted using crosses. - The transfer of electrons is shown using an arrow. Sodium Chloride Crystal Structure X-ray studies enable chemists to work out the arrangement of ions in a crystal. Sodium Chloride has a cubic structure. (1) The three-dimensional arrangement of ions is called a crystal lattice. Each sodium ion is surrounded by 6 chloride ions. Each chloride ion is surrounded by 6 sodium ions. How to Write the Formulas of Ionic Compounds A chemical formula is a way of representing a compound using symbols for the atoms present and numbers to show how many atoms of each element are present. - Ionic compounds are usually formed between the metals of groups I and II, and the non-metals of groups VI and VII. - Metals have a tendency to lose electrons and non-metals have a tendency to gain electrons. - An ionic compound is overall neutral. Therefore, there must be the same number of positive charges as negative charges present in the compound. Example – Write the formula for Aluminium Oxide Al = Group 3 so forms three positive charges O = Group VI so it forms two negative charges Need to find the lowest common multiple of 3 and 2 which is 6. We need to have 2 a +6 charge and a -6 charge so the compound is neutral overall. Al2O3 is the formula for Aluminium Oxide. Writing Formulas of Compounds with Group Ions It is not possible to predict the formulas of group ions from the periodic table so we must learn the most common ones off by heart. Name Formula Charge Hydroxide ion OH- -1 Nitrate ion NO3- -1 Hydrogencarbonate ion HCO3- -1 Permanganate ion MnO4- -1 Carbonate ion CO32- -2 Chromate ion CrO42- -2 Dichromate ion Cr2O72- -2 Sulfate ion SO42- -2 Sulfite ion SO32- -2 Thiosulfate ion S2O32- -2 Phosphate ion PO43- -3 Ammonium ion NH4+ +1 Write the formula for Sodium Sulfate Sodium ion = Na+ (1 positive charge) Sulfate ion = SO42- (2 negative charges) Need to have 2 positive charges and 2 negative charges Na2SO4 Write the formula for Calcium Hydrogencarbonate Calcium ion = Ca2+ Hydrogencarbonate ion =HCO3- Need to have two positive charges and two negative charges Ca(HCO3)2 Writing Formulas of Compounds Containing Transition Metals It is not possible to predict the charges of the ions of the d-block elements. d- block elements have variable valency (combining power). The main transition metals that show variable valency are: - Iron: Combines with chlorine to form either FeCl2 of FeCl3. FeCl2 is an ionic compound and the +2 charge on the iron atom is represented by putting the Roman numeral II in brackets. FeCl2 is called Iron (II) Chloride and FeCl3 is called Iron (III) Chloride. - Copper: Combines with oxygen to form either Cu2O – Copper (I) oxide, or CuO – Copper (II) oxide. - Chromium: Exists as Cr3+ ions in Chromium (III) chloride, CrCl3. 3 - Manganese – MnO2 is an ionic compound is called Manganese(IV) oxide, commonly called Manganese dioxide. Transition metals exhibit variable valency because there is such a small energy difference between the 4s and 3d sublevels. This means they can lose different numbers of electrons from these sublevels to give metal ions with different positive charges. Write the formula of Iron (II) carbonate. (II) indicates Fe2+. Carbonate ion = Co32-. Fe2+CO32- = Iron (II) Carbonate d-Block Elements and Transition Elements **Definition: A transition metal is one that forms at least one ion with a partially filled d sub-level.** Transition metals are the elements found in the d-block (except for Scandium and Zinc). They are special for several reasons: - Transition metals are used as catalysts. - Each transition metal can make different ions with different charges. For example, Mn (Manganese) can form ions with a charge of either +2, +4 or +7. We show which charge the ion has by using Roman Numerals in brackets after the name of the element. An Mn7+ ion is called Manganese (VII). - Transition metals form coloured compounds e.g., Mn7+ is purple, but Mn2+ is completely colourless. Covalent Bonding – Sharing of Electrons **Definition: A molecule is a group of atoms joined together. It is the smallest particle of an element or compound that can exist independently.** Covalent bonding involves electrons being shared between atoms, rather than electrons being fully transferred. A hydrogen molecule is formed when two individual hydrogen atoms share their singular electrons with each other. This forms a hydrogen molecule which is stable because it has two electrons in its outer energy level. 4 A chlorine atom has seven electrons in its outer shell. To create a stable octet, two chlorine atoms share a pair of electrons together forming a Chlorine molecule. The electrons are help together with a covalent bond. An oxygen atom has six electrons in the outer shell. This atom needs two more electrons to complete its octet. To achieve this, an oxygen atom shares two of its own electrons with the two hydrogen atoms. Two covalent bonds are formed. Two pairs of electrons on the oxygen are not involved in bonding, these are called lone pairs. Pairs of electrons that are involved in bonding are called bond pairs. Since a hydrogen atom cannot combine with more than one atom of any other element, its valency is known as monovalent. **Definition: The valency of an atom is defined as the number of atoms of hydrogen or any other monovalent element with which each atom of the element combines.** - Since chlorine combines with one atom of hydrogen, it has a valency of one. - Since oxygen combines with to atoms of hydrogen, it has a valency of two. Sigma and Pi bonds **Definitions: A sigma bond is formed by the head-on overlap of two orbitals. A pi bond is formed by the sideways overlap of p orbitals.** - A single bond always consists of a sigma bond. - A double bond always consists of one sigma and one pi bond. - A triple bond always consists of one sigma bond and two pi bonds. (2) 5 In the Nitrogen molecule we can see that the two 2px orbitals overlap head-on to form a sigma bond. The two 2py orbitals overlap sideways to form a pi bond. The two 2pz overlap sideways to form a second pi bond. Differing properties between Ionic and Covalent Compounds Ionic Covalent High melting and boiling points Low melting and boiling points Usually Solid at room temperature Usually liquid, gases, or soft solids at room temperature Usually hard and brittle Usually, soft Conduct electricity when molten or Do not conduct electricity dissolved in water Contain a network of ions in the Contain individual molecules crystal Shapes of Covalent Molecules In 1940, Sidgwick and Powell proposed a theory to account for the shapes of molecules. Their theory is known as the Valence Shell Electron Pair Repulsion Theory (VSEPR Theory). The shape of a molecule depends on the number of pairs of electrons around the central atom. As electrons are negatively charged, the electron pairs repel each other and arrange themselves in space so that they are as far apart as possible. 6 Method for using the VSEPR Theory - Draw a dot and cross diagram of the molecule - Count the bond pairs around the central atom - Count the lone pairs around the central atom - Use the number of bond pairs and lone pairs to work out the shape using the following table (Know off by heart) Shape Number of Number of Diagram Bond angles bond pairs lone pairs Linear 2 0 1800 V- Shaped 2 2 104.50 Triangular 3 0 1200 Planar Pyramidal 3 1 1070 Tetrahedral 4 0 109.50 Electronegativity – ‘Tug-Of-War’ for Electrons **Definition: Electronegativity is the relative attraction that an atom in a molecule has for the shared pair of electrons in a covalent bond.** - In a covalent bond between identical atoms, the pair of electrons is shared equally between the two atoms in the molecule. - Even though molecules have polar bonds, it doesn’t mean that the atom itself is polar - In bonds between different atoms, the pair of electrons is often more attracted to one of the atoms than the other. - This means that covalent molecules have two main types of bonds: - Non-polar Covalent bond: The atoms in the molecule all share electrons equally. **Definition: Polar Covalent – A polar covalent bond is a bond in which there is unequal sharing of the pair (or pairs) of electrons. This causes one end of the bond to be slightly positive (σ +) and the other end slightly negative (σ -). Uses of Electronegativity Values To Predict Polarity of Covalent Bonds By referring to the table of electronegativity values, we can tell if the bonds in a molecule are polar. The greater the electronegativity difference, the more polar the bond. 7 H2O HCl E.N of Hydrogen = 2.20 E.N of Hydrogen = 2.20 E.N of Oxygen = 3.44 E.N of Chlorine = 3.16 E.N difference = 1.24 E.N difference = 0.96 Highly polar covalent Polar covalent CH4 H-H E.N of Carbon = 2.55 E.N of Hydrogen = 2.20 E.N of Hydrogen = 2.20 E.N difference = 0 E.N difference = 0.35 Non-polar covalent Negligible polar covalent To Demonstrate whether a Liquid is Polar or Non-Polar (3) It can be proven that water is polar by placing a charged plastic rod near a thin stream of water. The stream of water is attracted to the negatively charged rod. This is because the positive poles of water molecules are attracted to the negatively charged rod. - If the rod is positively charged, the water molecules will spin so that the negative end of the water molecule is facing the rod, causing an attraction. - If the rod is negatively charged, the water molecules will spin so that the positive end of the molecule will face the rod, causing attraction - If the experiment is repeated with a non-polar liquid, no attraction is observed since there are no polar molecules in the liquid. Dissolving of Ionic Compounds in Water - Water is an excellent solvent because it is a polar molecule. Most ionic substances and most polar covalent substances dissolve in water. - An ionic substance such as sodium chloride will dissolve in water because the ionic bonding in NaCl is overcome by the strong attraction between the ions and the polar water molecules. - These ions are dragged away from the crystal lattice and become surrounded 8 by water molecules. (4) To Predict which Compounds are Ionic and which are Covalent - An electronegativity difference greater than 1.7 indicates ionic bonding in a compound. - An E.N difference less than or equal to 1.7 indicates covalent bonding in a compound. - An E.N difference greater than 0.4 and less than 1.7 indicates that the covalent bond is polar covalent. - An E.N difference less than or equal to 0.4 indicates that the covalent bond is non-polar. Intramolecular Bonding and Intermolecular Bonding **Definition: Intramolecular bonding is bonding that takes place within a molecule i.e., it holds the atoms together. Covalent bonding and polar covalent bonding are examples of intramolecular bonding.** **Definition: Intermolecular forces are the forces of attraction that exist between molecules. Van der Waals forces, dipole-dipole forces and hydrogen bonding are examples of intermolecular forces.** Van der Waals forces **Definition: Van der Waals forces are weak attractive forces between molecules resulting from the formation of temporary dipoles. They are the only forces of attraction between non-polar molecules.** - Small attractive forces which occur between non-polar molecules. - Imagine two electrons moving around inside a hydrogen molecule. As the two electrons move inside the molecule, it may happen that, at any one instant, both electrons may be closer to one end of the molecule than the other. - A temporary dipole is set up in the molecule which could induce a similar dipole in a nearby molecule. - There is then an attraction between the opposite charges, this is called a van der Waals force. - Van der Waals forces get stronger for larger molecules. Bigger molecules have more electrons so have a bigger chance of making a temporary dipole. - The stronger the Van der Waals forces on a molecule, the higher its boiling point and melting point. 9 Dipole-Dipole Forces **Definition: Dipole-Dipole forces are forces of attraction between the negative pole of one polar molecule and the positive pole of another polar molecule.** - Dipole-Dipole forces are very similar to Van der Waals forces, except that they occur between polar covalent molecules - Caused by the partially negative end of one molecule attracting the partially positive end of a neighbouring molecule. - Much weaker than the ionic bonds holding ionic compounds together but are much stronger than Van der Waals forces. - These forces cause polar molecules to have much higher boiling and melting points than similar non-polar molecules. Hydrogen Bonding **Definition: : Hydrogen Bonds are particular types of dipole-dipole attractions between molecules in which Hydrogen atoms are bonded to Nitrogen, Oxygen or Fluorine. The Hydrogen atom carries a partial positive charge and is attracted to the electronegative atom (N, O or F) in another molecule. Thus, the hydrogen bond acts as a bridge between two electronegative atoms in separate molecules.** - When Hydrogen is bonded to either Nitrogen, Oxygen or Fluoride, the partial charges are very strong as there is a large electronegativity difference between the atoms. This means that Hydrogen Bonding is the strongest form of intermolecular force. - Water has a high boiling point as the strong hydrogen forces need to be broken by a lot of heat for the liquid water to boil. 10 Exam Questions 2014 – HL – Section B – Question 4 (e) How many (i) sigma bonds, (ii) pi bonds, result from sharing of the valence electrons between the atoms in a molecule of nitrogen? 1 sigma and 2 pi 11. 11. Answer any two of the parts (a), (b) and (c). (2  25) (a) Define electronegativity. The relative attraction of an atom for a shared pair of electrons in a covalent bond Ammonia (NH3) and silane (SiH4) are small molecules, each of which has four electron pairs in the valence shell of the central atom. Account for the difference in bond angle between the two molecules, 107.3 in ammonia and 109.5 in silane. Lone pair of electron has greater repelling power than a bond pair of electrons. Ammonia has three bond pairs (one lone pair) where silane has four bond pairs Use electronegativity values to determine which bond, the N–H bond in ammonia or the Si–H bond in silane, is the more polar. Electronegativity difference greater for N – H therefore N – H more polar Which of the two substances has hydrogen bonding between its molecules? Ammonia Justify your answer. In ammonia hydrogen bonded to a small, highly electronegative element Give the reason why a molecule with polar bonds can be non-polar. Centres of positive and negative charge coincide 2013 – HL – Section B – Question 4 (d) Give the shape and the corresponding bond angle for a molecule of formula QX4 where Q is an element from Group 4 of the periodic table. Tetrahedral. 109.5° 10. Answer any two of the parts (a), (b) and (c) (a) Distinguish between intramolecular bonding and intermolecular forces Intramolecular: Forces between atoms in molecules. Intermolecular: Forces between molecules Explain each of the following in terms of intramolecular bonding or intermolecular forces or both. (i) The boiling point of hydrogen (20 K) is significantly lower than that of oxygen (90.2 K). Hydrogen is smaller, therefore weaker intermolecular forces (ii) Iodine has a very low solubility in water. Iodine is pure non-polar. Water is a polar solvent (iii) When a charged rod is held close to a thin stream of water flowing from a burette, the stream of water is deflected Charge on rod attracts opposite charge on polar water molecule 11 2012 – HL – Section B – Question 5 (d) Consider the following hydrides of some of the elements from the second and third periods of the periodic table: H2O NH3 PH3 HCl (i) State how the bonding in PH3 differs from the bonding in the other three hydrides. What is the reason for this difference in bonding? PH3 is virtually non- polar and other three are polar covalent. Tiny electronegativity difference in PH3 but much bigger electronegativity differences in the other three (ii) From these four hydrides, identify the hydride or hydrides in which hydrogen bonding occurs between the molecules. H2O and NH3 Give one property that is affected by the presence of intermolecular hydrogen bonding in the hydride or hydrides you have identified. Higher boiling point (iii) State the shape of the PH3 molecule and explain using electron-pair repulsion theory how this shape arises. Pyramidal. Repulsion between three bond pairs and the one lone pair. (e) Boron trichloride (BCl3) is a colourless gas. Would you expect (i) the B–Cl bonds, (ii) the BCl3 molecules, to be polar or non-polar? Justify your answers. B–Cl bond Polar BCI3 molecule: Non-polar Unequal sharing of electrons between B and CI cancels due to symmetry of molecule. 12  References 1. Shutterstock.com 2. Studysmarter.co.uk 3. Pdst.ie 4. Quora.com 13