5.1 b groups oxidation states
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This document discusses transition metals and their ionic charges. Transition metals are found in the middle of the periodic table and are called the B group elements. Unlike representative elements whose ionic charge is equal to their group number, transition metals can exist in multiple oxidation states by gaining or losing electrons. Examples given are iron (Fe2+ and Fe3+) and tungsten (W2+ through W7+). To predict the ionic charge of a transition metal, you must examine how it reacts with other elements and determine its oxidation state, which represents the charge of an atom in a compound. Rules for determining oxidation states are provided.
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This document provides an overview of ionic bonding, including:
1) Ionic bonds form between metals and nonmetals through the transfer of electrons, creating ions with opposite charges that are attracted to each other.
2) Cations are positively charged ions formed when metals lose electrons, while anions are negatively charged ions formed when nonmetals gain electrons.
3) Ionic compounds have crystalline structures and are strong conductors when molten or dissolved, due to the movement of ions.
oxidation-reduction reactions (redox reactions) a.pdf
oxidation-reduction reactions (redox reactions) are reactions in which electrons are
lost by an atom or ion in one reactant and gained by an atom or ion in another reactant. Although
electrons are gained and lost in these reactions, the balanced equation for a redox reaction does
not show the electrons that are being transferred. In order to tell whether a redox reaction has
occurred or not, we need a way to keep track of electrons. The best way to do so is by assigning
oxidation numbers to the atoms or ions involved in a chemical reaction. Oxidation numbers are
hypothetical numbers assigned to an individual atom or ion present in a substance using a set of
rules. Oxidation numbers (or oxidation states as they are also called) can be positive, negative, or
zero. It is VERY IMPORTANT to remember that oxidation numbers are always reported for one
individual atom or ion and not for groups of atoms or ions. The following rules are used to
assign oxidation numbers. Chem 1115 students will have these rules available on exams. Chem
1215 students must memorize these rules. Chem 1115 Chem 1215 Tutorial List Oxidation
Number Rules The oxidation number for an atom in its elemental form is always zero. A
substance is elemental if both of the following are true: only one kind of atom is present charge =
0 Examples: S8: The oxidation number of S = 0 Fe: The oxidation number of Fe = 0 The
oxidation number of a monoatomic ion = charge of the monatomic ion. Examples: Oxidation
number of S2- is -2. Oxidation number of Al3+ is +3. The oxidation number of all Group 1A
metals = +1 (unless elemental). The oxidation number of all Group 2A metals = +2 (unless
elemental). Hydrogen (H) has two possible oxidation numbers: +1 when bonded to a nonmetal -1
when bonded to a metal Oxygen (O) has two possilbe oxidation numbers: -1 in peroxides (O22-
)....pretty uncommon -2 in all other compounds...most common The oxidation number of
fluorine (F) is always -1. The sum of the oxidation numbers of all atoms (or ions) in a neutral
compound = 0. The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on
the polyatomic ion. When assigning oxidation numbers to the elements in a substance, take a
systematic approach. Ask yourself the following questions: Is the substance elemental? Is the
substance ionic? If the substance is ionic, are there any monoatomic ions present? Which
elements have specific rules? Which element(s) do(es) not have rules? Use rule 8 or 9 from
above to calculate these. Example: Determine the oxidation number of each element in Na2SO4.
Is the substance elemental? No, there are three elements present. Is the substance ionic? Yes,
because metal + non-metal = ionic. If the substance is ionic, are there any monoatomic ions
present? Yes, the sodium ion (Na+) is monoatomic. Therefore, the oxidation number of Na is +1.
Which elements have specific rules? Oxygen has a rule (-2 in most compounds). Oxidation
number of O = -2. Which element(s) do(es) not.
Electrochemistry
Includes a discussion of Voltaic and electrolytic cells, the Nernst equation and the relationship between electrochemical processes, chemical equilibrium and free energy.
**More good stuff available at:
www.wsautter.com
and
http://www.youtube.com/results?search_query=wnsautter&aq=f
Std XI-Ch-5-Redox-Reactions
This document discusses oxidation numbers, which indicate the degree of oxidation or reduction of atoms in chemical compounds. Oxidation numbers are assigned according to rules such as atoms in their elemental form having an oxidation number of 0, monoatomic ions having the same charge as their oxidation number, and the sum of oxidation numbers in a neutral molecule being 0. Oxidation numbers can be fractional and are indicated in Roman numerals after the symbol of a metal in its compound. An example of a fractional oxidation number is given for sulfur atoms in the tetrathionate ion.
Reaction intermediates
A reaction intermediate or an intermediate is a molecular entity that is formed from the reactants (or preceding intermediates) and reacts further to give the directly observed products of a chemical reaction.
1515753021_aqa-knowledge-mat-gcse-chemistry-p1.pptx
Rutherford's scattering experiment showed that the atom has a small, dense nucleus surrounded by electrons. Most alpha particles passed through the gold foil, but some were deflected or reflected, indicating a small, positively charged nucleus. The development of the atomic model over time led to the discovery of subatomic particles like protons, neutrons, and electrons. The periodic table arranges elements in order of atomic number and groups elements with similar properties together.
Oxidation reactions
Oxidation reactions in chemical engineering. Oxidation state. Oxidation state changes. Identify the element oxidized . Oxidation and reduction half-reactions.
Iron with hydrochloric acid . Zinc and copper. Aluminum and manganate. Cyanide and manganate. Production of ammonia from nitrite.
Balancing Oxidation Reduction Equations. The sulfite ion concentration present in wastewater from a papermaking plant.
Oxidizing and reducing agents
Lesson 1: Valence Electrons, Oxidation #, Dot Diagrams
This document provides an overview of physical science concepts related to subatomic particles, the periodic table, oxidation numbers, and Lewis dot diagrams. The key points are:
1) It reviews subatomic particles, periods and groups of the periodic table, and electron configurations.
2) The objective is to predict oxidation numbers and draw Lewis dot diagrams by understanding valence electrons.
3) It defines valence electrons, oxidation numbers, and ionic bonds in chemical compounds.
4) Examples are given of writing chemical symbols, determining valence electrons, and drawing Lewis dot diagrams using the cross method.
Metals, Non Metals And Oxidation
This document provides information about the periodic table, including the location and properties of metals, non-metals, and metalloids. It discusses periodic trends such as atomic radius and ionic charge. Various topics are covered, including oxidation-reduction reactions, ionic and covalent bonding, naming ionic and covalent compounds, and common polyatomic ions.
oxidation numbers
Oxidation numbers reflect the gain or loss of electrons by an atom in a compound. They are assigned according to a set of rules, such as hydrogen being +1, oxygen being -2, and the sum of oxidation numbers in a neutral compound being 0. Oxidation numbers track the change in electrons during oxidation-reduction reactions where oxidation is a loss of electrons and reduction is a gain.
Ch4&5_Chem60StudyGuide_W20.pdf
This document discusses key concepts in atomic structure and bonding:
1. It describes John Dalton's atomic theory and the discoveries of J.J. Thomson, Ernest Rutherford, and other scientists that led to developing the modern atomic model.
2. Key aspects of the modern atomic model are explained, including that atoms are composed of protons, neutrons, and electrons, with the nucleus at the center containing protons and neutrons.
3. The document also introduces concepts such as isotopes, ions, and how elements are arranged on the periodic table based on their atomic structure.
NOMENCLATURE OF BINARY POLY-ATOMIC COMPOUNDS
Giving names to chemical compounds, formulae, common polyatomic atoms. introduction about binary compounds, ionic bonds. provided with examples of binary ionic compounds.
A&P basic chemistry, atoms to ions, bonding, molecules v compounds, water and...
This document provides an overview of key concepts in chemistry including:
1) Matter is made up of atoms, which are the basic building blocks. Atoms can combine to form molecules or ionic compounds.
2) The periodic table describes all known elements and their properties. Atoms can gain or lose electrons to form ions and achieve stability.
3) Chemical bonds form when atoms share or transfer electrons. Ionic bonds form when an electron is transferred between atoms to form ions, while covalent bonds form when electrons are shared between atoms.
Similar to 5.1 b groups oxidation states (20)
oxidation-reduction reactions (redox reactions) a.pdf
oxidation-reduction reactions (redox reactions) a.pdf
1515753021_aqa-knowledge-mat-gcse-chemistry-p1.pptx
1515753021_aqa-knowledge-mat-gcse-chemistry-p1.pptx
Lesson 1: Valence Electrons, Oxidation #, Dot Diagrams
Lesson 1: Valence Electrons, Oxidation #, Dot Diagrams
A&P basic chemistry, atoms to ions, bonding, molecules v compounds, water and...
A&P basic chemistry, atoms to ions, bonding, molecules v compounds, water and...
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3.3 electron configuration
The document discusses electron configuration and energy levels. It introduces the historical models of the atom including Rutherford's plum pudding model and Bohr's model introducing energy levels. Modern models use orbitals that show the probability of finding electrons. Light can be described as both a wave and particle. Electrons absorb energy to move to excited states then emit light returning to ground states. Elements' electron configurations follow the aufbau principle filling lower energy orbitals first up to the periodic table.
3.3 orbital notation
Orbital notation represents the electron configuration of an element by showing the placement of each electron through arrows. It illustrates the Pauli exclusion principle, which states that no two electrons in the same atom can have the same quantum numbers, requiring opposite spin directions. The homework involves writing out electron configurations using orbital notation diagrams for various elements like sodium, boron, and titanium.
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This document discusses atomic theory and the key discoveries of Dalton. It explains that atoms are the fundamental building blocks of all matter, though they cannot be seen. Dalton proposed that elements are made of unique atom types, and compounds are made of two or more atom types combined in fixed ratios. The document reviews Dalton's five principles of atomic theory, including that atoms combine in whole number ratios. It provides examples of elemental and compound formulas showing atomic ratios.
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This document provides information for students taking Chemistry class with Mr. Martensen in the fall of 2012. It outlines classroom rules, grading policies, and requirements for a scientific calculator. Students must have their own scientific calculator, as they will not be provided. The class builds skills incrementally through units, so it is important not to fall behind. Mr. Martensen offers help sessions and maintains a class website to support students.
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The document contains a series of questions and answers about chemistry concepts like electron configurations, noble gas configurations, and atom to element ratios. It reviews elements like chlorine, sodium, iron, bromine, barium, neptunium, cobalt, silver, tellurium, radium, and ruthenium. It also addresses common errors in writing electron configurations and calculates atom to element ratios for compounds like H2O, H2SO4, C6H12O6, and KClO3.
Ch3 test prep forslideshare
The document contains examples of chemistry problems involving elements, electron configurations, and ratios of atoms to elements in compounds. It provides the answers to questions about protons, neutrons, and electrons in chlorine. It also gives electron configurations for various elements as well as which elements have specific noble gas configurations. Finally, it lists the atom to element ratios for common compounds like H2O, H2SO4, C6H12O6, and KClO3.
13.3 solubility
When ionic compounds dissolve in water, they break down into their constituent ions. Water is able to dissolve ionic compounds because it is a polar solvent. The solubility of a substance depends on temperature - solids generally become more soluble as temperature increases while gases become less soluble. A saturated solution contains as much solute as can dissolve at a given temperature, while an unsaturated solution contains less and a supersaturated solution contains more.
13.2 molarity
Molarity describes the concentration of a solution in terms of moles of solute per liter of solution. To calculate molarity, you need the moles of solute and the total volume of solution. The moles of solute can be determined by converting any given mass of solute to moles using the molar mass. The total volume must include both the solute and solvent volumes, as adding solute increases the total volume of the solution. Proper unit conversion and labeling of final answers is important when calculating molarity.
13.1 solutions
A solution is a homogeneous mixture of two or more substances uniformly dispersed throughout a single phase. A suspension is a heterogeneous mixture where particles remain mixed while being stirred but later settle to the bottom. A colloid is a middle-sized mixture that is stable and homogeneous like a solution but contains particles that can be seen with a light microscope.
12.3
The document discusses key concepts relating to gases:
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The document classifies different types of chemical reactions:
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Empirical formulas
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More from ZB Chemistry (20)
5.1 b groups oxidation states
- 1. Transition Metals & Ionic Charge Group B elements Courtesy of: Mr. Boroski
- 2. Representatives • We have talked about the Representative elements… • Groups 1A through 8A • And ion formation IonIon chargecharge GroupGroup numbernumber valencevalence numbernumber # of e-# of e- lost orlost or gainedgained
- 3. Transition elements • Found in the middle of the table • Known as the B group elements • Are called transition because they can change their charges! • How do they become charged? • By losing electrons.
- 4. Examples • Iron • Fe2+ • Fe3+ • Tungsten • W2+ • W3+ • W4+ • W5+ • W6+ • W7+
- 5. How do we predict??? • Cannot use group number! • Must see how it reacts with other elements. • Need to see its OxidationOxidation state (+ or -).state (+ or -).
- 6. Oxidation state • Represents the charge of an atom in a compound at a given time. – It is similar to the charge on the atom. • To figure, need to know the oxidation state rules…
- 7. OXIDATION RULES 1. Oxidation state of an element (in its pure form) is 0. 2. Oxidation state for representatives is their ionic charge most of the time. 3. Oxygen is always –2 unless in a peroxide, then it is –1. 4. H with nonmetals(covalent compounds) is given +1 oxidation state. Usually in the front of the formula. 5. The sum of oxidation numbers must be 0 for compounds and whatever the charge of the ionic species is.
- 8. To use rules: • You need to be able to do two things: An atom inventory Add and multiply whole numbers that are positive and negative.
- 9. Examples: Find oxidation states for all elements present. • CO2 C O 1 2 x x = = 0 -2 -4 +4+4 Total charge of compound #ofatoms times Oxidationstate Totalchargeperelement Elementsymbol equals
- 10. Examples: Find oxidation states for all elements present. • CuCl2 Cu Cl 1 2 x x = = 0 -1 -2 +2+2
- 11. Examples: Find oxidation states for all elements present. • H2SO4 H S 2 1 x x = = 0 +6 +6 +2+1 O 4 x =-2 -8
- 12. Examples: Find oxidation states for all elements present. • NO3 1- N O 1 3 x x = = -1 -2 -6 +5+5
- 13. Homework: Oxidation numbers problems. • DIRECTIONS: find the oxidation numbers of ALL elements in the compounds. 1. CuCl 6. ZnBr2 2. CuI2 7. PbNO3 3. Fe2 O3 8. Mg3 N2 4. Ag3 N 9. KCl 5. KMnO4 10. CrO4 -2