How To Calculate Oxidation Numbers - Basic Introduction

The Organic Chemistry Tutor

22 Aug 201731:15

EducationalLearning

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TLDRThis video script provides an in-depth guide to understanding oxidation numbers, starting with the basics such as the oxidation state of pure elements and ions. It then delves into the rules for determining oxidation states in compounds, emphasizing the role of electronegativity. The script uses numerous examples to illustrate how to calculate the oxidation state of elements in various compounds and polyatomic ions, highlighting the relationship between electronegativity values and typical oxidation states. The video concludes with a discussion on average oxidation states in compounds with multiple oxidation states for a single element.

Takeaways

📌 The oxidation state of a pure element is always zero, including elements like zinc, oxygen, fluorine, and phosphorus.
🔋 The oxidation state of ions is equivalent to their charge, such as Zn^2+ having a +2 oxidation state and Fe^3+ having a +3 oxidation state.
🤝 For diatomic ions, divide the total charge by the number of atoms to find the individual oxidation states, like with Hg^2+2 where each mercury has a +1 oxidation state.
📚 In compounds, fluorine typically has a -1 oxidation state, while oxygen usually has a -2 oxidation state, except in peroxides (-1) and superoxides (-1/2).
🔄 Hydrogen has a +1 oxidation state when bonded to non-metals and a -1 oxidation state when bonded to metals, based on its relative electronegativity.
🔢 To find oxidation states in compounds, use the rule that the algebraic sum of the oxidation states of all atoms must equal zero.
📈 Electronegativity increases across a period from left to right on the periodic table, with fluorine being the most electronegative element.
🔍 Use electronegativity values to predict the oxidation states of elements in compounds, especially when they form bonds with different elements.
🌐 In compounds containing polyatomic ions, calculate the oxidation state of each element based on the overall charge of the ion.
📊 Average oxidation states can be decimal values when different ions within a compound have different oxidation states, such as in Fe3O4 where iron can have +2 or +3 oxidation states.
🌟 Remember that individual ions in a compound must have whole number oxidation states, even if their average is a decimal, indicating a mix of different oxidation states among the ions.

Q & A

What is the oxidation state of a pure element?
-The oxidation state of a pure element is always zero, as there are no charges involved and it is not a compound.
How do you determine the oxidation state of ions?
-The oxidation state of an ion is equivalent to the charge of the ion. For example, a zinc ion with a 2+ charge has an oxidation state of +2.
What is the oxidation state of fluorine when it is part of a compound?
-Fluorine always has an oxidation state of -1 when it is part of a compound, as it is the most electronegative element.
How does electronegativity influence oxidation states?
-Electronegativity is the ability of an atom to attract electrons to itself. The element with higher electronegativity typically has a negative oxidation state, while the less electronegative element usually has a positive oxidation state.
What is the oxidation state of oxygen in the peroxide ion?
-In the peroxide ion, each oxygen atom has an oxidation state of -1, as the total charge of the ion is divided by the number of oxygen atoms to find the individual oxidation states.
What is the oxidation state of magnesium in the compound MgCl2?
-Magnesium has an oxidation state of +2 in the compound MgCl2, balancing the -1 oxidation state of each chlorine atom.
How do you find the oxidation state of elements in polyatomic ions?
-To find the oxidation state of elements in polyatomic ions, you set up an equation where the sum of the oxidation states of all atoms equals the net charge of the ion. Then solve for the unknown oxidation state.
What is the average oxidation state of iron in Fe3O4?
-The average oxidation state of iron in Fe3O4 is 2.67. This indicates that in the compound, there are two iron ions with a +3 charge and one with a +2 charge, to maintain electrical neutrality.
How does the oxidation state of an element change when it is bonded to metals versus non-metals?
-When an element is bonded to a metal, it typically has a negative oxidation state because metals are less electronegative. Conversely, when bonded to non-metals, the element often has a positive oxidation state due to the non-metals being more electronegative.
What is the oxidation state of sulfur in H2SO4 (sulfuric acid)?
-In H2SO4, sulfur has an oxidation state of +6. This is determined by considering the oxidation states of hydrogen (+1) and oxygen (-2) and solving for the unknown sulfur oxidation state to achieve a net charge of zero for the molecule.
What is the relationship between the oxidation state and the number of electrons an atom can attract?
-The oxidation state is directly related to the number of electrons an atom can attract. A higher oxidation state indicates a greater number of electrons attracted by the atom, typically because it has a higher electronegativity.

Outlines

00:00

📚 Introduction to Oxidation Numbers

This paragraph introduces the concept of oxidation numbers, explaining how to determine the oxidation state of elements. It emphasizes the rule that pure elements have an oxidation state of zero and delves into the oxidation states of ions, including monoatomic and diatomic ions. The paragraph also discusses the importance of understanding electronegativity in determining oxidation states, especially when dealing with compounds.

05:00

🔍 Understanding Electronegativity and Oxidation States

The second paragraph focuses on the relationship between electronegativity and oxidation states. It provides examples of how the most electronegative element in a compound typically carries a negative charge, while less electronegative elements carry a positive charge. The paragraph also explains how to calculate the oxidation states of elements in compounds and polyatomic ions, using examples involving vanadium, oxygen, sulfur, and phosphorus.

10:00

🌟 Electronegativity Values and Oxidation States

This paragraph provides specific electronegativity values for common elements and uses these values to determine the oxidation states of oxygen and fluorine in oxygen difluoride. It explains how the most electronegative element will attract electrons more strongly, resulting in a partial negative charge, while the less electronegative element will have a partial positive charge. The paragraph also applies this concept to solve for the oxidation states in hydrochloric acid and sodium hydride.

15:04

📊 Calculating Oxidation States with Electronegativity

The fourth paragraph continues to explore the concept of electronegativity in determining oxidation states. It provides a method for calculating the oxidation states of elements in compounds like sulfuric acid, sulfur dioxide, ammonia, and nitrous oxide. The paragraph emphasizes the use of periodic tables and the understanding of electronegativity to deduce the typical oxidation states of elements.

20:07

🔢 Dealing with Average Oxidation States

This paragraph discusses the concept of average oxidation states, which can result in non-whole numbers. It uses examples of propane and iron(III) oxide to illustrate how to calculate the average oxidation state when a compound contains multiple ions with different oxidation states. The explanation clarifies that while the average oxidation state can be a decimal, individual ions will have whole number charges.

25:10

🌐 Polyatomic Ions and Mixed Oxidation States

The final paragraph of the script focuses on the oxidation states within polyatomic ions that contain multiple different elements. It provides examples of how to determine the oxidation states of each element in complex ions, such as the sulfate ion and the bicarbonate ion. The paragraph concludes with examples involving bromine and chlorine, and bromine and iodine, highlighting the variable oxidation states of transition metals and the influence of electronegativity on the oxidation states of elements in compounds.

Mindmap

Keywords

💡Oxidation Numbers

Oxidation numbers, also known as oxidation states, are used to keep track of electrons when dealing with redox reactions. They represent the charge an atom would have if all bonds were ionic. In the video, the concept is introduced as a method to determine the charge of an element within a compound or ion, with examples like zinc having an oxidation state of +2 when it forms an ion.

💡Electronegativity

Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It increases across a period on the periodic table and is used to predict the oxidation state of elements in compounds. In the video, it is emphasized that the more electronegative element typically has a negative oxidation state, while the less electronegative element has a positive oxidation state.

💡Pure Elements

Pure elements refer to substances that are composed of only one type of atom. In the context of oxidation numbers, the oxidation state of a pure element is always zero since there are no charges involved. The video emphasizes this rule when discussing how to find oxidation states.

💡Ions

Ions are atoms or molecules that have a net electrical charge due to the loss or gain of one or more electrons. The video explains that the oxidation state of an ion is equal to its charge. It also discusses monoatomic and diatomic ions, and how to calculate the oxidation state of elements in polyatomic ions.

💡Compounds

Compounds are substances formed when two or more chemical elements are chemically bonded together. The video explains how to determine the oxidation states of elements within compounds by using known oxidation states of common elements like fluorine and oxygen, as well as the concept of electronegativity.

💡Polyatomic Ions

Polyatomic ions are charged groups of atoms that act as a single ion in chemical reactions. The video explains how to find the oxidation states of elements within polyatomic ions by considering the overall charge of the ion and the known oxidation states of its constituent atoms.

💡Redox Reactions

Redox reactions, short for reduction-oxidation reactions, involve the transfer of electrons between two species. Oxidation states are crucial in these reactions as they help to keep track of the electrons transferred. The video does not explicitly mention redox reactions but focuses on the concept of oxidation numbers which is fundamental to understanding redox processes.

💡Charge Conservation

Charge conservation is a principle stating that the total charge in a chemical reaction must remain constant. In the context of the video, this principle is used to determine the oxidation states of elements in compounds and ions, ensuring that the sum of the oxidation states of all atoms in a neutral compound is zero.

💡Alkaline Earth Metals

Alkaline earth metals are a group of elements in the second column of the periodic table, including elements like magnesium and calcium. These metals typically have two electrons in their outermost shell and tend to form cations with a +2 oxidation state when they lose these electrons.

💡Halogens

Halogens are elements in the seventeenth group of the periodic table, including fluorine, chlorine, bromine, iodine, and astatine. They are highly electronegative and typically form anions with a -1 oxidation state when they gain an electron.

💡Transition Metals

Transition metals are elements in the middle of the periodic table and are known for their ability to form ions with multiple oxidation states. They are characterized by their ability to easily lose or gain electrons and are often involved in redox reactions.

Highlights

The oxidation state of a pure element is always zero.

For ions, the oxidation state is the same as the ion's charge.

In diatomic ions, the oxidation state can be found by dividing the total charge by the number of atoms.

In compounds, fluorine typically has a -1 oxidation state.

Oxygen usually has a -2 oxidation state in compounds, except in peroxides (-1) and superoxides (-1/2).

Hydrogen has a +1 oxidation state when bonded to non-metals and -1 when bonded to metals.

The more electronegative element typically carries a negative charge.

The oxidation state of magnesium in MgCl2 is +2, balancing the -1 charge of chlorine.

In AlF3, aluminum has a +3 oxidation state, and fluorine maintains its -1 state.

In vanadium oxide (VO2), vanadium has a +5 oxidation state, and oxygen has a -2 state.

The sulfate ion (SO4) has sulfur in a +6 oxidation state.

In phosphate (PO4), phosphorus has a +5 oxidation state.

Nitrogen in nitrate (NO3) has a +5 oxidation state, while chlorine in perchlorate (ClO4) has a +7 oxidation state.

Electronegativity increases towards fluorine on the periodic table, affecting oxidation states.

In oxygen difluoride (OF2), oxygen has a +2 oxidation state due to fluorine's higher electronegativity.

In hydrochloric acid (HCl), chlorine has a -1 oxidation state, and hydrogen has a +1 state.

Sodium hydride (NaH) has sodium with a +1 and hydrogen with a -1 oxidation state.

In BH3, boron has a +3 oxidation state because hydrogen is more electronegative.

In sulfuric acid (H2SO4) and sulfur dioxide (SO2), the oxidation states of sulfur and oxygen differ based on electronegativity.

In NH3, nitrogen has a -3 oxidation state, and in NO2, it has a +4 state due to different electronegativity interactions.

In methane (CH4), carbon has a -4 oxidation state, and in carbon dioxide (CO2), it has a +4 state.

Average oxidation states can be decimal values, as seen in C3H8 and Fe3O4, indicating a mix of different oxidation states for individual atoms.

Polyatomic ions with three different elements, such as in K2CrO4, require solving for the transition metal's oxidation state based on electronegativity and charge balance.

In potassium bicarbonate (KHCO3), carbon has a +4 oxidation state, determined by breaking down the compound into its ions.

In BrCl3 and IBr5, the oxidation states are determined by the relative electronegativities of bromine and chlorine, and iodine, respectively.

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