Page 1: Content Information

• Document Details

  • Title: QUÍMICA INORGÁNICA

  • Author: Lic. Mario Santiago Zetina

  • Year: 2012

Page 2: Basic Equations

Key Formulas

1. Media Velocity Formula:

   • v = d/t

   • Rearranging for t:

   • t = d/v

2. Density Formula:

   • ρ = m/V

   • Rearranging for m:

   • m = ρ \cdot V

3. Ohm's Law:

   • V = I \cdot R

   • Rearranging for I:

   • I = V/R

4. Kinetic Energy Formula:

   • E_k = 1/2 \cdot mv²

   • Rearranging for v:

   • v = √(2E_k/m)

Page 3: More Equations

Additional Formulas

5. Force Formula:

   • F = m \cdot a

   • Rearranging for a:

   • a = F/m

6. Circle Area Formula:

   • A = πr²

   • Rearranging for r:

   • r = √(A/π)

7. Specific Heat Capacity:

   • Q = m \cdot c \cdot ΔT

   • Rearranging for c:

   • c = Q/(m \cdot ΔT)

8. Universal Gravity Equation:

   • F = G(m₁m₂/r²)

   • Rearranging for r:

   • r = √(G(m₁m₂)/F)

Page 4: References

• Textbook Reference:

  • Title: QUÍMICA

  • Edition: 13ª edición

  • Authors: Raymond Chang and Jason Overby

  • Publisher: McGraw Hill

Page 5: History of Chemistry

Historical Milestones

• Prehistory: Discovery of fire, pottery, and paint production.

• Antiquity: Greek philosophers adopt the concept of matter.

• Alchemy: Search for the philosopher's stone and the elixir of life.

Page 6: Key Figures and Developments

Significant Contributions

• Iatrochemistry: Medicinal chemistry.

• Antoine Lavoisier (1776): Known as the "father of chemistry"; detailed explanation of combustion and scientific method development.

• Modern Chemistry: Discovery of chemical elements, atomic theory, organic chemistry, etc.

Page 7: What is Chemistry?

Definition and Scope

• Chemistry is the study of matter and its changes.

• It deals with substances and their interactions.

• Focuses on preparation, properties, and transformations of substances.

Page 8: Branches of Chemistry

Major Branches

• Pure Chemistry: Studies all types of substances, organic and inorganic.

• Organic Chemistry: Focused on carbon compounds.

• Inorganic Chemistry: Studies structure, composition, and reactions of elements and compounds.

• Analytical Chemistry: Analyzes properties of chemical systems.

• Industrial Chemistry: Examines production methods of chemical substances.

Page 9: Basic Concepts

Definitions

• Matter: Anything that occupies space and has mass (visible and invisible).

• Substance: A form of matter with a defined composition and distinct properties (e.g., water, sugar).

• Mixture: Combination of two or more substances retaining their properties (e.g., milk, cement).

Page 10: Types of Mixtures

Classification of Mixtures

• Homogeneous Mixtures: Uniform composition (elements indistinguishable).

• Heterogeneous Mixtures: Non-uniform composition (distinguishable substances).

Page 11: Examples of Homogeneous Mixtures

Visual Representations

1. Example: Saltwater is a homogeneous mixture.

2. Visual Representation: Figure illustrating uniformity of saltwater solution.

Page 12: Examples of Heterogeneous Mixtures

Visual Representations

1. Visual Representation: Image illustrating distinct phases of heterogeneous mixtures.

Page 13: Identifying Mixtures

Examples and Questions

1. Gold and silver alloy (homogeneous).

2. Activated carbon in a liquid (heterogeneous).

3. Carbon dioxide in water (homogeneous).

4. Silica and clay in water (heterogeneous).

5. Soil with small stones (heterogeneous).

Page 14: Separation of Mixtures

Key Concept

• Mixtures can be formed and separated through physical means without changing the identity of components (e.g., boiling saltwater).

Page 15: Elements vs. Compounds

Definitions

• Element: A substance that cannot be separated into simpler substances by chemical means.

• Compound: A substance formed from two or more elements chemically bonded in fixed proportions.

  • Separation: Compounds only separable by chemical methods.

Page 16: Common Elements and Symbols

Element List

• Aluminum (Al), Arsenic (As), Sulfur (S), Barium (Ba), Bismuth (Bi), Bromine (Br).

Page 17: Diagrams of Elements and Compounds

Representation

• Visual representations of atoms indicating their classification as elements or compounds.

Page 18: Classification of Matter

Hierarchical Structure

• Matter can be separated into:

  • Mixtures

    • Homogeneous

    • Heterogeneous

  • Substances

    • Compounds

    • Elements

Page 19: States of Matter

Three States

• Solid

• Liquid

• Gas

Page 20: Characteristics of States of Matter

Properties

• Solid: Atoms/molecules are closely packed, minimal movement.

• Liquid: Atoms/molecules close but not rigid; can move.

• Gas: Atoms/molecules are spread apart; large distances between them.

Page 21: Questions about States of Matter

Inquiry Questions

• Why are gases more compressible than liquids or solids?

• What happens molecularly when solids change to liquids (melting)?

• Why do solids have defined shapes while liquids and gases do not?

Page 22: Phase Changes

Key Points

• Melting Point: Temperature where solid becomes liquid.

• Boiling Point: Temperature where liquid becomes gas.

• Condensation Point: Temperature where liquid becomes solid.

Page 23: Physical and Chemical Properties

Types of Properties

• Physical Property: Can be observed/measured without changing substance (color, melting point).

• Chemical Property: Alters molecular structure during interactions (e.g., hydrogen combustion).

Page 24: Chemical Change

Concept Overview

• Reaction that leads to the formation of products with different properties from the reactants.

Page 25: Examples of Properties

Property Identification

1. Iron forming rust (chemical).

2. Metal stretching into wire (physical).

3. Temperature at which liquid becomes gas (physical).

4. Copper reacts with sulfuric acid (chemical).

5. Vinegar and baking soda reaction (chemical).

6. Wood burning (chemical).

7. Reflective surface (physical).

8. Cooling reduces a substance's volume (physical).

Page 26: Extensive vs. Intensive Properties

Definitions

• Extensive Properties: Depends on amount of substance (e.g., mass, volume).

• Intensive Properties: Independent of amount of substance (e.g., temperature, density).

Page 27: Properties Comparison

Summary of Properties

• Intensive: Temperature, boiling point, density.

• Extensive: Weight, size, volume.

Page 28: Diagrams of Elements

Visual Comparison

• Diagrams portraying compounds vs. physical/chemical changes.

Page 29: Notation

Data Presentation

• Notation information placeholder (context needed).

Page 30: Measurements

Types of Properties

• Macroscopic Properties: Can be directly measured.

• Microscopic Properties: Cannot be directly observed.

SI Units

1. Time

2. Mass

3. Volume

4. Density

5. Temperature

Page 31: Measurement Tools

Common Instruments

• Measuring flask, graduated cylinder, burette, pipette.

Page 32: SI Units Overview

Basic Units

• Length: Meter (m)

• Mass: Kilogram (kg)

• Time: Second (s)

• Electric Current: Ampere (A)

• Temperature: Kelvin (K)

• Amount of Substance: Mole (mol)

• Luminous Intensity: Candela (cd)

Page 33: Mass vs. Weight

Key Differences

• Mass: Amount of matter in an object.

• Weight: Gravity's force on an object.

SI Units

• Mass: kg, Volume: m³.

Page 34: Density Exercises

Practical Exercise Example

• Cylinder A: 800 g, 1000 cm³, Density = 0.8 g/cm³

• Cylinder B: 1000 g, 2000 cm³, Density = 0.5 g/cm³

• Conclusion: Cylinder A has higher density.

Page 35: Gold Properties

Example Calculation

• Gold sample: 301 g, Volume: 15.6 cm³

• Density = 19.3 g/cm³.

Page 36: Density Practice Problems

Exercises

1. Platinum density problem.

2. Mercury density problem.

Page 37: Temperature Scales

Key Temperature Points

• Fahrenheit: Freezing 32°F, Boiling 212°F.

• Celsius: Freezing 0°C, Boiling 100°C.

• Kelvin: Absolute temperature scale.

Page 38: Temperature Conversion Exercises

Sample Conversion Problems

1. Convert Celsius to Fahrenheit (80 °C).

2. Convert Fahrenheit to Celsius (130 °F).

Page 39: Scientific Notation

Overview

• Scientific notation representation details for calculations.

Page 40: Scientific Notation Application

Importance

• Used to simplify multiplication/division of large/small numbers in chemistry.

Page 41: Finding N in Scientific Notation

Explanation

• Description of moving decimal for scientific notation.

Page 42: Special Cases in Notation

Concept Overview

• Handle special cases in scientific notation.

• n = 0 implication.

Page 43: Arithmetic with Scientific Notation

Examples of Operations

1. Addition, subtraction, multiplication of scientific notation formats.

Page 44: Significant Figures

Representation of Figures

• Different notations for significant figures

Page 45: Definition of Significant Figures

Explanation

• The last digit in a measurement is always uncertain, indicating precision of measurements.

Page 46: Rules for Significant Figures

Specific Guidelines

1. Non-zero digits are significant.

2. Captive zeros are significant.

3. Leading zeros are not significant.

4. Trailing zeros are significant if there's a decimal present.

Page 47: Significant Figure Problems

Practice

• Determine significant figures given various numbers.

Page 48: Rules for Calculations

Operations with Significant Figures

1. Addition/Subtraction: Depends on decimal places.

2. Multiplication/Division: Depends on significant figures.

Page 49: Additional Calculation Guidelines

Redundant Rules

• Specifics on how to handle rounding in significant figures during addition and subtraction.

Page 50: Accuracy vs. Precision

Definitions

• Accuracy: Closeness to true value.

• Precision: Closeness among multiple measurements.

Page 51: Example Scenario

Measurement Task

• Comparing students’ mass measurements for accuracy and precision.

Page 52: Precision and Accuracy Test

Example Problem

• Measurement task showcasing accuracy and precision in science.

Page 53: Conceptual Test

Understanding Inquiry

• Evaluating performance in terms of accuracy or precision.

Page 54: Measurement Data Comparison

Data Presentation

• Players’ scores representing precision and accuracy comparisons.

Page 55: Measuring Metal Mass

Measurement Results

• Series of mass measurements for a metal cube and evaluation of accuracy and precision.

Page 56: Conversion Factors

Key Concept

• Conversion factors facilitate transitioning between measurement units.

Page 57: Practical Application of Conversion Factors

Example Exercises

1. How many cm in 5.84 m?

2. Convert kg to lb.

3. Glucose daily intake conversion.

4. Volume conversion for helium.

5. Density of mercury in kg/m³.

Page 58: Additional Conversion Exercises

Sample Problems

1. Convert gallons to cubic centimeters.

2. Density to volume.

Page 59: Final Conversion Exercises

Measurement Practices Example

1. Convert gallons to cubic centimeters.

2. Temperature and density calculations.

Page 60: Thermochemistry

Fundamental Concepts

• Energy capacity in work production.

Page 61: Types of Energy

Energy Classifications

• Kinetic, radiant, thermal, chemical, and potential energy.

Page 62: Energy in Reactions

Energy Transfer

• All chemical reactions absorb or release energy, often as heat.

Page 63: Additional Concepts

Placeholder Content

Page 64: Thermodynamic Systems

Systems Classifications

• Open, closed, and isolated systems defined by energy and mass transfer.

Page 65: Exothermic vs. Endothermic Reactions

Energy Changes

• Definitions and examples of energy release and absorption in reactions.

Page 66: Application of Energy Concepts

Understanding Systems

• Classifying systems and reactions based on energy characteristics.

Page 67: System Evaluation

Problem-Solving

• Classify system types and determine reaction types for given scenarios.

Page 68: First Law of Thermodynamics

Core Principles

• Energy cannot be created or destroyed, only transformed.

Page 69: Calorimetry

Overview of Measurements

• Calorimetry, specific heat, and capacity heat definitions and formulas explained.

Page 70: Specific Heat Calculation Example

Sample Calculation

• Calculation of heat based on specific heat values.

Page 71: Heat Transfer Calculations

Equations for Heat Calculation

• Detailed equations describing heat absorption/release.

Page 72: Heat Absorption Example

Practical Calculation Problem

• Example calculating heat absorbed by water during temperature increase.

Page 73: Copper Heat Calculation

Heat Calculation Scenario

• Example calculating heat necessary to change the temperature of copper.

Page 74: Heat Release Calculation Problem

Example Heat Calculation

• Evaluating heat released during a cooling process.

Page 75: Mass Calculation Example

Assessing Heat Capacity

• Problem determining mass based on specific heat and temperature change.

Page 76: Initial Temperature Determination

Problem-Solving Guide

• Task determining the initial temperature of a heated aluminum sheet.

Page 77: Short Exam Example

Test Questions

• Set of exercises demonstrating measurement and conversion skills in chemistry.

Page 78: Short Exam Instructions

Exam Content

• Guidelines and examples for answering chemistry measurement problems.