Undervisningsbeskrivelse
Stamoplysninger til brug ved prøver til gymnasiale uddannelser
|
Termin(er)
|
2024/25
|
|
Institution
|
X - Ikast-Brande Gymnasium
|
|
Fag og niveau
|
Kemi -
|
|
Lærer(e)
|
Lene Malmgaard
|
|
Hold
|
2024 ChemHL (1i ChemHL)
|
Oversigt over gennemførte undervisningsforløb
Beskrivelse af de enkelte undervisningsforløb (1 skema for hvert forløb)
|
Titel
1
|
Topic 1- Atomic structure
SL+HL
Structure 1.1.1—Elements are the primary constituents of matter, which cannot be chemically broken down into simpler substances.
Compounds consist of atoms of different elements chemically bonded together in a fixed ratio.
Mixtures contain more than one element or compound in no fixed ratio, which are not chemicallynbonded and so can be separated by physical methods.
Distinguish between the properties of elements, compounds and mixtures.
Structure 1.1.2—The kinetic molecular theory is a model to explain physical properties of matter (solids, liquids and gases) and changes of state.
Distinguish the different states of matter.
Use state symbols (s, l , g and aq) in chemical equations.
Structure 1.1.3—The temperature, T, in Kelvin (K) is a measure of average kinetic energy Ek of particles.
Interpret observable changes in physical properties and temperature during changes of state.
Convert between values in the Celsius and Kelvin scales.
Structure 1.2.1—Atoms contain a positively charged, dense nucleus composed of protons and neutrons (nucleons). Negatively charged electrons occupy the space outside the nucleus.
Use the nuclear symbol AZ X to deduce the number of protons, neutrons and electrons in atoms and ions.
Structure 1.2.2—Isotopes are atoms of the same element with different numbers of neutrons.
Perform calculations involving non-integer relative atomic masses and abundance of isotopes from givendata.
Structure 1.3.1—Emission spectra are produced by atoms emitting photons when electrons in excited states return to lower energy levels.
Qualitatively describe the relationship between colour, wavelength, frequency and energy across the electromagnetic spectrum.
Distinguish between a continuous and a line spectrum.
Structure 1.3.2—The line emission spectrum of hydrogen provides evidence for the existence of electrons in discrete energy levels, which converge at higher energies.
Describe the emission spectrum of the hydrogen atom, including the relationships between the lines and energy transitions to the first, second and third energy levels.
Structure 1.3.3—The main energy level is given an integer number, n, and can hold a maximum of 2n2 electrons.
Structure 1.3.4—A more detailed model of the atom describes the division of the main energy level into s, p, d and f sublevels of successively higher energies.
Structure 1.3.5—Each orbital has a defined energy state for a given electron configuration and chemical environment, and can hold two electrons of opposite spin.
Sublevels contain a fixed number of orbitals, regions of space where there is a high probability of finding an electron.
Apply the Aufbau principle, Hund’s rule and the Pauli exclusion principle to deduce electron
configurations for atoms and ions up to Z = 36
Structure 3.1.1—The periodic table consists of periods, groups and blocks.
Identify the positions of metals, metalloids and non-metals in the periodic table.
Structure 3.1.2—The period number shows the outer energy level that is occupied by electrons.
Elements in a group have a common number of valence electrons.
Deduce the electron configuration of an atom up to Z = 36 from the element’s position in the periodic table and vice versa.
HL only
Structure 1.2.3—Mass spectra are used to determine the relative atomic masses of elements from their isotopic composition.
Interpret mass spectra in terms of identity and relative abundance of isotopes.
Structure 1.3.6—In an emission spectrum, the limit of convergence at higher frequency corresponds to ionization.
Explain the trends and discontinuities in first ionization energy (IE) across a period and down a group.
Calculate the value of the first IE from spectral data that gives the wavelength or frequency of the convergence limit.
Structure 1.3.7—Successive ionization energy (IE) data for an element give information about its electron configuration.
Deduce the group of an element from its successive ionization data.
Practicals: Paper chromatography of plant pigments or sweets. Flame colours (demo).
TOK:
- Do atomic orbitals exist or are they primarily useful inventions to aid our understanding?
- Which of Daltons five proposals do we hold to be true today? How does scientific knowledge change with time? Are models and theories accurate descriptions or just useful interpreations that help to explain the natural world?
ATL:
Thinking skills: being curious and openminded to new ideas in science.
Social skills: working in groups in the lab
Self management skills: Being punctual and meeting deadlines
|
|
Indhold
|
Kernestof:
Skriftligt arbejde:
| Titel |
Afleveringsdato |
|
Assignment 1 - Significant figures
|
18-09-2024
|
|
Assignment 1 Significant figures
|
18-09-2024
|
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
13 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
|
Titel
2
|
Topic 2 Stoichiometry - mole calculations
Structure 1.4.1—The mole (mol) is the SI unit of amount of substance. One mole contains exactly the number of elementary entities given by the Avogadro constant.
Convert the amount of substance, n, to the number of specified elementary entities.
Structure 1.4.2—Masses of atoms are compared on a scale relative to 12C and are expressed as relative atomic mass Ar and relative formula mass Mr .
Determine relative formula masses Mr from relative atomic masses Ar .
Structure 1.4.3—Molar mass M has the units g mol–1.
Solve problems involving the relationships between the number of particles, the amount of substance in moles and the mass in grams.
Structure 1.4.4—The empirical formula of a compound gives the simplest ratio of atoms of each element present in that compound. The molecular formula gives the actual number of atoms of each element present in a molecule.
Interconvert the percentage composition by mass and the empirical formula.
Determine the molecular formula of a compound from its empirical formula and molar mass.
Structure 1.4.5—The molar concentration is determined by the amount of solute and the volume of solution.
Solve problems involving the molar concentration, amount of solute and volume of solution.
Structure 1.4.6—Avogadro’s law states that equal volumes of all gases measured under the same conditions of temperature and pressure contain equal numbers of molecules.
Solve problems involving the mole ratio of reactants and/or products and the volume of gases.
Structure 1.5.1—An ideal gas consists of moving particles with negligible volume and no
intermolecular forces. All collisions between particles are considered elastic.
Recognize the key assumptions in the ideal gas model.
Structure 1.5.2—Real gases deviate from the ideal gas model, particularly at low temperature and high pressure.
Explain the limitations of the ideal gas model.
Structure 1.5.3—The molar volume of an ideal gas is a constant at a specific temperature and pressure.
Investigate the relationship between temperature, pressure and volume for a fixed mass of an ideal gas and analyse graphs relating these variables.
Structure 1.5.4—The relationship between the pressure, volume, temperature and amount of an ideal gas is shown in the ideal gas equation PV = nRT and the combined gas law P1V1/
T1 = P2V2/ T2.
Solve problems relating to the ideal gas equation.
Reactivity 2.1.1—Chemical equations show the ratio of reactants and products in a reaction.
Deduce chemical equations when reactants and products are specified
Reactivity 2.1.2—The mole ratio of an equation can be used to determine:
• the masses and/or volumes of reactants and products
• the concentrations of reactants and products for reactions occurring in solution.
Calculate reacting masses and/or volumes and concentrations of reactants and products.
Reactivity 2.1.3—The limiting reactant determines the theoretical yield.
Identify the limiting and excess reactants from given data.
Reactivity 2.1.4—The percentage yield is calculated from the ratio of experimental yield to
theoretical yield.
Solve problems involving reacting quantities, limiting and excess reactants, theoretical, experimental and percentage yields.
Reactivity 2.1.5—The atom economy is a measure of efficiency in green chemistry.
Calculate the atom economy from the stoichiometry of a reaction.
Practicals:
Molar mass of lighter gas
Empirical formula of Magnesium Oxide
Titration of HCl with NaOH
Uncertainties in measurements, how to propagate them and assess their importance.
TOK:
- The magnitude of Avogadros number is beyond the scale of our everyday experience. How do we perceive and imagine concepts of the microscopic world?
- The ideal gas equation can be deduced based on a list of assumptions and general physics. What is the role of perception, imagination and intuition i development of scientific models?
ATL:
Communication in assignments:
• Using terminology, symbols and communication conventions consistently
and correctly
• Presenting data appropriately
Grade talks: self management skill
• Setting learning goals and adjusting them in response to experience
• Seeking and acting on feedback
|
|
Indhold
|
Kernestof:
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
19 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
|
Titel
3
|
Topic 3 Bonding
Structure 2.1.1—When metal atoms lose electrons, they form positive ions called cations.
When non-metal atoms gain electrons, they form negative ions called anions.
Predict the charge of an ion from the electron configuration of the atom.'
Structure 2.1.2—The ionic bond is formed by electrostatic attractions between oppositely charged ions. Deduce the formula and name of an ionic compound from its component ions, including polyatomic ions. Binary ionic compounds are named with the cation first, followed by the anion. The anion adopts the suffix “ide”.
Interconvert names and formulas of binary ionic compounds.
Structure 2.1.3—Ionic compounds exist as three-dimensional lattice structures, represented by empirical formulas.
Explain the physical properties of ionic compounds to include volatility, electrical conductivity and solubility.
Structure 2.2.1—A covalent bond is formed by the electrostatic attraction between a shared pair of electrons and the positively charged nuclei.The octet rule refers to the tendency of atoms to gain a valence shell with a total of 8 electrons.
Deduce the Lewis formula of molecules and ions for up to four electron pairs on each atom.
Structure 2.2.2—Single, double and triple bonds involve one, two and three shared pairs of
electrons respectively.
Explain the relationship between the number of bonds, bond length and bond strength.
Structure 2.2.3—A coordination bond is a covalent bond in which both the electrons of the shared pair originate from the same atom.
Identify coordination bonds in compounds.
Structure 2.2.4—The valence shell electron pair repulsion (VSEPR) model enables the shapes of molecules to be predicted from the repulsion of electron domains around a central atom.
Predict the electron domain geometry and the molecular geometry for species with up to four electron domains.
Structure 2.2.5—Bond polarity results from the difference in electronegativities of the bonded atoms.
Deduce the polar nature of a covalent bond from electronegativity values.
Structure 2.2.6—Molecular polarity depends on both bond polarity and molecular geometry.
Deduce the net dipole moment of a molecule or ion by considering bond polarity and molecular geometry.
Structure 2.2.7—Carbon and silicon form covalent network structures.
Describe the structures and explain the properties of silicon, silicon dioxide and carbon’s allotropes: diamond, graphite, fullerenes and graphene.
Structure 2.2.8—The nature of the force that exists between molecules is determined by the size and polarity of the molecules. Intermolecular forces include London (dispersion), dipole-induced dipole, dipole–dipole and hydrogen bonding.
Deduce the types of intermolecular force present from the structural features of covalent molecules.
Structure 2.2.9—Given comparable molar mass, the relative strengths of intermolecular forces are generally: London (dispersion) forces < dipole–dipole forces < hydrogen bonding.
Explain the physical properties of covalent substances to include volatility, electrical conductivity and solubility in terms of their structure.
Structure 2.2.10—Chromatography is a technique used to separate the components of a mixture based on their relative attractions involving intermolecular forces to mobile and stationary phases.
Explain, calculate and interpret the retardation factor values, RF.
Structure 2.3.1—A metallic bond is the electrostatic attraction between a lattice of cations and delocalized electrons.
Explain the electrical conductivity, thermal conductivity and malleability of metals.
Structure 2.3.2—The strength of a metallic bond depends on the charge of the ions and the radius of the metal ion.
Explain trends in melting points of s and p block metals.
Structure 2.4.1—Bonding is best described as a continuum between the ionic, covalent and metallic models, and can be represented by a bonding triangle.
Use bonding models to explain the properties of a material
Structure 2.4.2—The position of a compound in the bonding triangle is determined by the relative contributions of the three bonding types to the overall bond.
Determine the position of a compound in the bonding triangle from electronegativity data.
Predict the properties of a compound based on its position in the bonding triangle.
Structure 2.4.3—Alloys are mixtures of a metal and other metals or non-metals. They have
enhanced properties.
Explain the properties of alloys in terms of non-directional bonding.
Structure 2.4.4—Polymers are large molecules, or macromolecules, made from repeating subunits called monomers.
Describe the common properties of plastics in terms of their structure.
Structure 2.4.5—Addition polymers form by the breaking of a double bond in each monomer.
Represent the repeating unit of an addition polymer from given monomer structures.
HL only:
Structure 2.2.11—Resonance structures occur when there is more than one possible position for a double bond in a molecule.
Deduce resonance structures of molecules and ions.
Include the term “delocalization”.
Structure 2.2.12—Benzene, C6H6, is an important example of a molecule that has resonance.
Discuss the structure of benzene from physical and chemical evidence.
Structure 2.2.13—Some atoms can form molecules in which they have an expanded octet of
electrons. Visually represent Lewis formulas for species with five and six electron domains around the central atom.
Deduce the electron domain geometry and the molecular geometry for these species using the VSEPR model.
Structure 2.2.14—Formal charge values can be calculated for each atom in a species and used to determine which of several possible Lewis formulas is preferred.
Apply formal charge to determine a preferred Lewis formula from different Lewis formulas for a species.
Structure 2.2.15—Sigma bonds σ form by the head-on combination of atomic orbitals where the electron density is concentrated along the bond axis.
Pi bonds π form by the lateral combination of p-orbitals where the electron density is
concentrated on opposite sides of the bond axis.
Deduce the presence of sigma bonds and pi bonds in molecules and ions.
Structure 2.2.16—Hybridization is the concept of mixing atomic orbitals to form new hybrid
orbitals for bonding.
Analyse the hybridization and bond formation in molecules and ions.
Identify the relationships between Lewis formulas, electron domains, molecular geometry and type of hybridization.
Predict the geometry around an atom from its hybridization, and vice versa.
Structure 2.3.3—Transition elements have delocalized d-electrons.
Explain the high melting point and electrical conductivity of transition elements
Structure 2.4.6—Condensation polymers form by the reaction between functional groups in each monomer with the release of a small molecule.
Represent the repeating unit of polyamides and polyesters from given monomer structures.
Practicals:
- Solubility of ionic compounds
- Making bioplastics
Project: The plastic problem, group presentations.
TOK/NOS:
- To what extent does collaboration and competition help or hinder production of knowledge? Use idea of large plastic molecules as example.
- The bonding triangle is a tool with predictive power about propertives of substances
(NOS)
- Electron orbitals and hybridization does not exast as physical entities, but provides a convinient model for explaining electron properties and bonding. To what extent does language and imagery enhance or mask understanding of what is represented?
ATL:
Collaboration: Working collaboratively to achieve a common goal
• Assigning and accepting specific roles during group activities
• Appreciating the diverse talents and needs of others
Communication: Clearly communicating complex ideas in response to open-ended
questions
• Using digital media for communicating information
Thinking skills:
Experimenting with new strategies for learning
|
|
Indhold
|
Kernestof:
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
30 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
|
Titel
4
|
Topic 4 Kinetics - rate of reaction
Reactivity 2.2.1—The rate of reaction is expressed as the change in concentration of a particular reactant/product per unit time.
Determine rates of reaction
Reactivity 2.2.2—Species react as a result of collisions of sufficient energy and proper orientation.
Explain the relationship between the kinetic energy of the particles and the temperature in kelvin, aqnd the role of collision geometry
Reactivity 2.2.3—Factors that influence the rate of a reaction include pressure, concentration, surface area, temperature and the presence of a catalyst.
Predict and explain the effects of changing conditions on the rate of a reaction
Reactivity 2.2.4—Activation energy, Ea, is the minimum energy that colliding particles need for a successful collision leading to a reaction.
Construct Maxwell–Boltzmann energy distribution curves to explain the effect of temperature on the probability of successful collisions.
Reactivity 2.2.5—Catalysts increase the rate of reaction by providing an alternative reaction
pathway with lower Ea.
Sketch and explain energy profiles with and without catalysts for endothermic and exothermic reactions.
Construct Maxwell–Boltzmann energy distribution curves to explain the effect of different values for Ea on the probability of successful collisions.
HL only:
Reactivity 2.2.6—Many reactions occur in a series of elementary steps. The slowest step determines the rate of the reaction.
Evaluate proposed reaction mechanisms and recognize reaction intermediates.
Distinguish between intermediates and transition states, and recognize both in energy profiles of reactions.
Reactivity 2.2.7—Energy profiles can be used to show the activation energy and transition state of the rate-determining step in a multistep reaction.
Construct and interpret energy profiles from kinetic data.
Reactivity 2.2.8—The molecularity of an elementary step is the number of reacting particles taking part in that step.
Interpret the terms “unimolecular”, “bimolecular” and “termolecular”.
Reactivity 2.2.9—Rate equations depend on the mechanism of the reaction and can only be
determined experimentally.
Deduce the rate equation for a reaction from experimental data.
Reactivity 2.2.10—The order of a reaction with respect to a reactant is the exponent to which the concentration of the reactant is raised in the rate equation.
The order with respect to a reactant can describe the number of particles taking part in the rate-determining step.
The overall reaction order is the sum of the orders with respect to each reactant.
Sketch, identify and analyse graphical representations of zero, first and second order reactions.
Reactivity 2.2.11—The rate constant, k, is temperature dependent and its units are determined from the overall order of the reaction.
Solve problems involving the rate equation, including the units of k.
Reactivity 2.2.12—The Arrhenius equation uses the temperature dependence of the rate constant to determine the activation energy.
Describe the qualitative relationship between temperature and the rate constant.
Analyse graphical representations of the Arrhenius equation, including its linear form.
Reactivity 2.2.13—The Arrhenius factor, A, takes into account the frequency of collisions with proper orientations.
Determine the activation energy and the Arrhenius factor from experimental data.
Practicals:
Effect of concentration on rate of reaction ( thiosulfate and acid).
Factors affecting rate of reaction (Design practical, each group decide on which factor to explore)
TOK:
Is there a fundamental differnece between knowledge claims based on theoretical data and experimental data? rate of reaction can only be determined experimentally.
Are physical properties such as temerpature invedted or discovered? we have two temperature scales Celcius and Kelvin, Celcius is based on properties of water whereas kelvin is based on average kinetic energy of gas molecules.
ATLs:
• Designing procedures and models
• Reflecting on the credibility of results
• Providing a reasoned argument to support conclusions
• Using terminology, symbols and communication conventions consistently
and correctly
• Presenting data appropriately
|
|
Indhold
|
Kernestof:
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
17 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
|
Titel
5
|
Topic 5 Equlibrium
Reactivity 2.3.1—A state of dynamic equilibrium is reached in a closed system when the rates of forward and backward reactions are equal.
Describe the characteristics of a physical and chemical system at equilibrium.
Reactivity 2.3.2—The equilibrium law describes how the equilibrium constant, K, can be
determined from the stoichiometry of a reaction.
Deduce the equilibrium constant expression from an equation for a homogeneous reaction.
Reactivity 2.3.3—The magnitude of the equilibrium constant indicates the extent of a reaction at equilibrium and is temperature dependent.
Determine the relationships between K values for reactions that are the reverse of each other at the same temperature.
Reactivity 2.3.4—Le Châtelier’s principle enables the prediction of the qualitative effects of changes in concentration, temperature and pressure to a system at equilibrium.
Apply Le Châtelier’s principle to predict and explain responses to changes of systems at equilibrium
HL only:
Reactivity 2.3.5—The reaction quotient, Q, is calculated using the equilibrium expression with non-equilibrium concentrations of reactants and products.
Calculate the reaction quotient Q from the concentrations of reactants and products at a particular time, and determine the direction in which the reaction will proceed to reach equilibrium
.
Reactivity 2.3.6—The equilibrium law is the basis for quantifying the composition of an equilibrium mixture.
Solve problems involving values of K and initial and equilibrium concentrations of the components of an equilibrium mixture
Reactivity 2.3.7—The equilibrium constant and Gibbs energy change, ΔG, can both be used to measure the position of an equilibrium reaction.
Practicals:
Factors affecting position of equilibrium, Le Chaletier in practice.
TOK:
What is the role of imagination and intuition in the creation of hypothesis in natural science? dynamic equilibria involves both the macroscopic and the microscopic scale.
What is the difference between theories and analogies as forms of explanation?
ATL:
• Combining different ideas in order to create new understandings
• Applying key ideas and facts in new contexts
• Engaging with, and designing, linking questions
• Experimenting with new strategies for learning
• Reflecting at all stages of the assessment and learning cycles
|
|
Indhold
|
Kernestof:
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
11 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
|
Titel
6
|
Topic 6 Proton transfer reactions- acids and bases
Reactivity 3.1.1—Brønsted–Lowry acid is a proton donor and a Brønsted–Lowry base is a proton acceptor.
Deduce the Brønsted–Lowry acid and base in a reaction.
Reactivity 3.1.2—A pair of species differing by a single proton is called a conjugate acid–base pair.
Deduce the formula of the conjugate acid or base of any Brønsted–Lowry base or acid.
Reactivity 3.1.3—Some species can act as both Brønsted–Lowry acids and bases.
Interpret and formulate equations to show acid–base reactions of these species.
Reactivity 3.1.4—The pH scale can be used to describe the [H+] of a solution:
pH = –log10[H+]; [H+] = 10–pH
Perform calculations involving the logarithmic relationship between pH and [H+].
Reactivity 3.1.5—The ion product constant of water, Kw, shows an inverse relationship between [H+] and [OH–]. Kw = [H+] [OH–]
Recognize solutions as acidic, neutral and basic from the relative values of [H+] and [OH–].
Reactivity 3.1.6—Strong and weak acids and bases differ in the extent of ionization.
Recognize that acid–base equilibria lie in the direction of the weaker conjugate.
Reactivity 3.1.7—Acids react with bases in neutralization reactions.
Formulate equations for the reactions between acids and metal oxides, metal hydroxides,
hydrogencarbonates and carbonates.
Reactivity 3.1.8—pH curves for neutralization reactions involving strong acids and bases have characteristic shapes and features.
Sketch and interpret the general shape of the pH curve.
HL only:
Reactivity 3.1.9—The pOH scale describes the [OH–] of a solution. pOH = –log10[OH–]; [OH–] = 10–pOH
Interconvert [H+], [OH–], pH and pOH values.
The equations for pOH are given in the data booklet.
Reactivity 3.1.10—The strengths of weak acids and bases are described by their Ka, Kb, pKa or pKb values.
Interpret the relative strengths of acids and bases from these data.
Reactivity 3.1.11—For a conjugate acid–base pair, the relationship Ka × Kb = Kw can be derived from the expressions for Ka and Kb.
Solve problems involving these values.
Reactivity 3.1.12—The pH of a salt solution depends on the relative strengths of the parent acid and base.
Construct equations for the hydrolysis of ions in a salt, and predict the effect of each ion on the pH of the salt solution.
Reactivity 3.1.13—pH curves of different combinations of strong and weak monoprotic acids and bases have characteristic shapes and features.
Interpret the general shapes of pH curves for all four combinations of strong and weak acids and bases.
Reactivity 3.1.14—Acid–base indicators are weak acids, where the components of the conjugate acid–base pair have different colours. The pH of the end point of an indicator, where it changes colour, approximately corresponds to its pKa value.
Construct equilibria expressions to show why the colour of an indicator changes with pH.
Reactivity 3.1.15—An appropriate indicator for a titration has an end point range that coincides with the pH at the equivalence point.
Identify an appropriate indicator for a titration from the identity of the salt and the pH range of the indicator
Reactivity 3.1.16—A buffer solution is one that resists change in pH on the addition of small
amounts of acid or alkali.
Describe the composition of acidic and basic buffers and explain their actions.
Reactivity 3.1.17—The pH of a buffer solution depends on both:
• the pKa or pKb of its acid or base
• the ratio of the concentration of acid or base to the concentration of the conjugate base or
acid.
Solve problems involving the composition and pH of a buffer solution, using the equilibrium constant.
Practicals:
- Reactions of acids with metal, carbonates and alkalis.
- pH titration curves using digitial pH sensors for collecting data (Technology skill)
- pH measurements of salt solutions
TOK:
In science we often make assumptions in scince to simplify mathematical calculations, also in calculation of pH in solutions of weak acids. What kind of justification is needed for that?
pH of neutral solutions is not always 7, that is surprising to many students. To what degree does prior knowledge hinder new understandings?
ATL:
• Reflecting on the credibility of results
• Providing a reasoned argument to support conclusions
• Combining different ideas in order to create new understandings
• Applying key ideas and facts in new contexts
• Appreciating the diverse talents and needs of others
• Resolving conflicts during collaborative work
• Actively seeking and considering the perspective of others
• Reflecting on the impact of personal behaviour or comments on others
|
|
Indhold
|
Kernestof:
|
|
Omfang
|
Estimeret:
Ikke angivet
Dækker over:
16 moduler
|
|
Særlige fokuspunkter
|
|
|
Væsentligste arbejdsformer
|
|
{
"S": "/lectio/1223/stamdata/stamdata_edit_student.aspx?id=666\u0026prevurl=studieplan%2fuvb_hold_off.aspx%3fholdid%3d80011813079",
"T": "/lectio/1223/stamdata/stamdata_edit_teacher.aspx?teacherid=666\u0026prevurl=studieplan%2fuvb_hold_off.aspx%3fholdid%3d80011813079",
"H": "/lectio/1223/stamdata/stamdata_edit_hold.aspx?id=666\u0026prevurl=studieplan%2fuvb_hold_off.aspx%3fholdid%3d80011813079"
}