Chemical Bonding: Ionic Bonding PDF

Chemical Bonding: Ionic Bonding First of all, What is a chemical bond? ➔ A chemical bond forms when outer electrons of atoms interact and rearrange into a more stable configuration. ➔ All chemical bonds rely on electrostatic attraction between positive and negative electrons...

Chemical Bonding: Ionic Bonding First of all, What is a chemical bond? ➔ A chemical bond forms when outer electrons of atoms interact and rearrange into a more stable configuration. ➔ All chemical bonds rely on electrostatic attraction between positive and negative electrons ◆ Electrostatic: describes the attraction between particles with opposite charges—like positively charged protons and negatively charged electrons—which helps atoms bond together. Ionic Bonding ➔ Occurs when metal atoms combine with non-metal atoms ◆ Example: Sodium Chloride Metals: Alkali Metals, Alkaline Earth Metals, Transition metals, Post Transition Metals, Non Metals: Hydrogen, Non-metals, Halogens, Noble gasses Ionic Bonding mechanism/ reasoning: ➔ Metal atoms have a low number of valence electrons so it is easier for them to lose them so they can gain a full octet. ➔ Non-metals have many valence electrons so they tend to gain electrons to reach a stable octet. Metals in Groups 1, 2, and 13 have 1, 2, or 3 electrons in their outer shell, respectively. To reach a stable configuration, they tend to lose these electrons (Cations): Group 1 metals lose 1 electron, forming ions with a 1+ charge. Group 2 metals lose 2 electrons, forming ions with a 2+ charge. Group 13 metals lose 3 electrons, forming ions with a 3+ charge. Non-metals in Groups 15, 16, and 17 have 5, 6, or 7 electrons in their outer shell, respectively. To complete their octet, they gain electrons (Anions): Group 15 elements gain 3 electrons, forming ions with a 3− charge. Group 16 elements gain 2 electrons, forming ions with a 2− charge. Group 17 elements gain 1 electron, forming ions with a 1− charge. Transition metals also lose electrons when forming ionic compounds; however their more complex electron configuration means that they can generally form more than one type of ion. For example, iron can form ions with either a 2+ or a 3+ charge. RULES FOR NAMING: - the positive ion is generally written first, followed by the negative ion - Eg: NaCl= Sodium Chloride (Sodium loses one electrons making it a cation) (Chlorine gain one electron making it an anion) - As compounds do not carry an overall charge, it is necessary to balance the charges of the anion and cation components - Eg: HNO3 = Hydrogen Nitrate Covalent Bonds A covalent bond is a chemical in which it shares electron pairs between atoms. This bond is also known as a molecular bond. Molecules must undergo covalent bonding How is this bond developed? A molecule is formed when a group of atoms are bonded together by covalent compounds. The attraction between the atomic nuclei and the shared electrons balances the repulsive forces between the nuclei, keeping the atoms at a stable distance from one another. - SINGLE BOND (one pair of electrons, 2 in total = DOUBLE BOND (two pairs of electrons, 4 in total = TRIPLE BOND (three pairs of electrons, 6 in total Examples: CH4 -Basically what you’re seeing in that image is the depiction of H’s only valence electron being shared with the main element C. O2 -In this image, there are 2 of the same elements given to create a stable configuration. As O has 7 valence electrons, the repeated elements share a double bond How to identify what element is sharing and receiving - Know that the element with the least electronegativity can be the central atom  - The elements on the right side of the periodic table tend to acquire electrons on the outermost shell to follow the octet rule - While the elements on the left side of the periodic table tend to give away their valence electrons What is Metallic Bonding? - Metal atoms in their elemental state have loosely held outer electrons. - These outer electrons tend to become delocalized (free to move). - Metal atoms lose these electrons and become positively charged ions (cations). - Cations form a lattice structure through which the delocalized electrons can move freely. - The attraction between this lattice of cations and the free electrons is called metallic bonding. The greater the number of delocalized electrons and the smaller the cation, the greater the binding force between them. The strength of the metallic bond is determined by: 1. The number of delocalized electrons 2. The charge of the cation 3. The radius of the cation Strength VS melting point The strength of the metallic bond tends to decrease down a group as the size of the cation increases, reducing the attraction between the delocalized electrons and the positive charges. Mg Na K Rb Melting Point 650 98 63 39 o / C *Transition Elements tend to have very high metallic bonds due to the large number of electrons that can become delocalized. Metal Physical Properties Good electrical conductivity: Delocalized electrons move easily through the metal in response to voltage. Good thermal conductivity: Delocalized electrons and tightly packed ions allow efficient heat transfer. Malleable and ductile: Non-directional movement of electrons keeps the metallic bond intact even when the metal shape changes under pressure. High melting points: Strong metallic bonds require a lot of energy to break and separate atoms. Shiny, lustrous appearance: Delocalized electrons in the metal's crystal structure reflect light. Ductility the measure of how readily a material can be turnt into a wire. One ounce can be 50 miles long of a wire (GOLD AGAIN) Malleable the ability of a substance to be hammered/rolled into thin thin sheets (GOLD Alloys ❖ Made by adding one metal (or carbon) to another in the molten state. ❖ As the mixture solidifies, ions of different metals scatter through the lattice, bound by delocalized electrons, forming metallic bonds. ❖ Alloys have unique properties due to different packing of cations in the lattice. ❖ Alloys are generally more chemically stable, stronger, and more resistant to corrosion than their component elements.